Slide 1

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 2

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 3

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 4

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 5

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 6

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 7

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 8

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 9

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 10

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 11

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 12

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 13

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 14

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 15

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 16

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 17

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 18

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 19

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 20

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 21

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 22

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 23

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 24

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 25

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 26

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 27

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 28

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 29

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 30

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 31

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 32

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 33

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?


Slide 34

Ionic Bonds
Chapter 4

Ionic Compounds


What are
Chemical
Bonds






Force that holds 2 atoms
together
Attraction between + nucleus and
– electron
Attraction between + ion and –
ion
Valence electrons make bonds

Ionic compounds


What
are
chemical
bonds?
(cont)



Elements react to form a stable
octet (noble gas configuration).






The + and – charges act like opposite
poles of a magnet.
Opposites attract strongly
Likes repel strongly
Strength diminishes with distance

Typical Ions


Oxidation number = oxidation state =
number of electrons transferred from an
atom to make a compound




Na+ oxidation number = +1
O2- oxidation number = 2Used to determine compound formulas

Ionic Compounds


How are
positive ions
formed?





Atom loses one or more VALENCE
electrons
Called a CATION
Ion becomes more stable by losing
electrons (octet rule)







Not a change in atom, Just an ion
Loses all electrons in outer shell

Reactivity depends on ease of losing
electrons
Transition metals usually form 2+ or
3+ ions shown with a (II) or (III)

Ionic Bonds


How do
negative ions
form?



Atoms gain negative electrons








Nonmetals have a great attraction for
electrons
Adding electron fills up the shell = stable

Called an ANION
Naming: change name to end in –ide
Gaining enough electrons to fill outer
shell (octet rule)




7A gains 1
6A gains 2
5A gains 3

Ionic Compounds


What are
oxidation
states



Oxidation state is the charge of the
‘normal’ ion formed









Group 1 loses 1 valence electron (+1)
Group 2 loses 2 valence electrons (+2)
Group 13 loses 3 valence electrons (+3)
Group 14 does not generally make ionic
compounds
Group 17 gains 1 valence electron (-1)
Group 16 gains 2 valence electrons (-2)
Group 15 gains 3 valence electrons (-3)

Ionic Compounds
Ionic compounds are Metal+ and
Nonmetal Metals make Cations








Nonmetals make Anions




Groups 1A (1) – 3A (13) and all Group D
elements
Form + ions
Group 5A (15) - 7A (17)

Nobel Gasses (Group 8A/18) do not form
compounds. Why?

Properties of Ionic Compounds


Crystal shape





High melting points




Alternate positive and negative ions in patterns
Crystals are all the same shape for each
compound
Table Salt Melting point is 801oC

Conduct electricity




Dissolve in water
Ions become more loosely associated
Pass electrical charges along

Counting Atoms in a Compound
Element Symbol

Element Symbol

CH4

Subscript

Formula Names




Remember: total number of electrons lost by
cations must equal total number of electrons
gained by anions!
Metal name is stated first.




Number of cations in ratio is subcripted

Nonmetal name is stated second



Suffix –ide is used
Number of anions in ration is subscripted

Formula Names


Examples


Sodium (Na+) and Chlorine (Cl-)
NaCl = ratio 1:1
Sodium Chloride



Calcium (Ca2+) and Fluorine (Fl-)

CaFl2 = ratio 1:2
Calcium Fluoride
Aluminum (3+) and Sulfur (2-)
Find lowest common dominator
(6)
2(Al 3+) + 3 (S 2+) = both transfer 6 electrons
Al2S3
Aluminum Sulfide

Determining Formulas from Oxidation
States
1.

Find oxidation numbers for elements


2.

Put elements together with oxidation numbers
as superscripts


3.

Ca+2O-2

Criss-Cross the oxidation numbers removing the
signs


4.

Ca, O

Ca+2O-2  Ca2O2

Take lowest common denominator


CaO

Formulas with Polyatomic Ions



Polyatomic ions stay together as a single
group
Example Calcium and Phosphate







Calcium (2+)
Phosphate PO4 (3-)
3 (Ca) + 2 ( PO4)
Ca3(PO4)2
Calcium Phosphate

Covalent Bonds
Chapter 9

Electron Sharing





Covalent bonds are nonmetal/nonmetal
Covalent (Co = together, Valence
Electrons)
Covalent Bonds – Formed when two atoms
SHARE a PAIR of electrons
Both atoms attract the shared electron(s) at
the same time

CoValent Bonds
Every atom MUST have
a full valence electron
shell:
Carbon = 8
Hydrogen = 2

Shared electrons count for
BOTH atoms

Covalent Bonds
 Some

compounds share more
than one pair of electrons

Molecular Compounds




Molecular Compounds have covalent
bonds
Much lower boiling points than ionic
Poor conductors of electricity





Pure water does not conduct electricity
Water with sugar does not conduct electricity
Water with SALT DOES conduct electricity

Naming Covalent
Molecules
Chapter 9.2

Basic Rules





First element in formula is always named
first, using the entire element name
Second element in formula is named
second, using the root of the element name
and adding the suffix –ide
Prefixes are added to each name to
indicate number of atoms of each type.

Prefixes












1 = mono2 = di3 = tri4 = tetra5 = penta6 = hexa7 = hepta8 = octa9 = nona10 = deca

Example


H2O (water)


Hydrogen – Oxygen






Oxide

DIhydrogen oxide

NH3 (ammonia)






Nitrogen – Hydrogen
1 nitrogen, 3 hydrogens
Nitrogen, 3 HydrIDE
Nitrogen TRI hydride

Metallic Bonds and
Metals
Chapter 4

Metallic Bonds


In solid state,m Metals do not bond ionically
but for form lattices




Similar to ionic crystal lattices

Metals have at least one valence electron
but:




do not share these electrons
Do not lose electrons

Metallic Bonds


In the solid crystal lattice of metals





Electrons are crowded
Outer energy levels overlap
Like "an array of positive ions in a sea of electrons".
Electrons are delocalized


Can move from one atom to another easily

Metallic Bonds


Metallic Bond


Attraction of a metallic cation for a delocalized
electron

Properties of Metals



Melting points vary greatly
Metals are







malleable – can be hammered into sheets
Ductile – can be drawn into wires
Generally durable – but with some variation
Good conductors – due to delocalized electrons

Mobile electrons consist of:



‘d’-level electrons
2 outer ‘s’ electrons

Metals
Malleability and ductility
Metals are described as malleable (can be beaten into
sheets) and ductile (can be pulled out into wires). This is
because of the ability of the atoms to roll over each other into
new positions without breaking the metallic bond.
If a small stress is put onto the metal, the layers of atoms will
start to roll over each other. If the stress is released again,
they will fall back to their original positions. Under these
circumstances, the metal is said to be elastic.

Metals
If a larger stress is put on, the atoms roll over each
other into a new position, and the metal is
permanently changed.

Metal Alloys


Easy to introduce other metals into the
metallic crystal




Properties of alloys differ some from either
base metal




Mixture called an alloy

Steel is iron mixed with another element

Alloys form when:



Elements involved are similar in size
Or one is much smaller than the other

Metal Alloys


2 types of alloys





Substitutional
Interstitial

Sustitutional alloys


Original metallic atoms are replaced by other
metal atoms of similar size


Sterling silver (copper and silver), brass, pewter,
10-carat gold

Metal Alloys


Interstitial alloys


Small holes in metal crystal are filled with other,
smaller atoms






Like pouring sand in a bucket of gravel
Example: carbon steel (iron crystal is filled with
carbon)
Physical properties of steel are changed. Iron is
relatively soft and brittle. Adding Carbon makes
the solid harder, stronger less ductile, giving it
different uses.

Making Ionic Compounds


Write the electron configuration of Magnesium.




Write the electron configuration for oxygen








Based on this configuration, will oxygen gain or lose
electrons?

When we burn magnesium, What compound results?




Based on this configuration, will Mg gain or lose electrons?

Magnesium + oxygen

Which atom donated electrons? How Many?
What is the formula for the new compound?
Do you think this new compound will conduct
electricity?