Transcript Introduction to Bonding Essential Question: What are the characteristics

Introduction to Bonding
Essential Question: What are the characteristics
of an ionic bond and how are they formed?
What is a bond?
• A bond can be thought of as a force that
holds groups of two or more atoms
together and makes them function as a
single unit
• Example : water
O
H
H
Bonds require energy to break
and release energy when made
Types of bonds
• Ionic bonds - typically formed between
metals and nonmetals
• Covalent bonds - typically formed
between nonmetals
• Metallic bonds - formed between metals
Ionic Bonds
• Ionic bonding results from the transfer of electrons.
Then, the opposite charges attract each other.
• Ionic bonds are strong
• Ion: Atoms that have lost or gained electrons and
have gained a charge
• Oxidation number: Charge of the ion (positive or
negative)
Ionic Bonds
• Metals lose electrons and become positive
(cations)
• Nonmetals gain electrons and become negative
(anions)
• Either way, atoms become ions and gain a
charge
Predicting Ionic Charges
1+
3+
1- 0
234+/-
Ionic Bonds
• Na and Cl
– Na is a metal and likes to lose one electron
and form a +1 ion.
– Cl is a nonmetal and likes to gain one electron
and form a -1 ion.
– the final ionic compounds is NaCl
Na+ + Cl-
NaCl
The electrostatic interaction
keeps them together!
Ionic Bonds
• They do this
to achieve an
octet!
Atomic Stability
1
Noble gases have a
filled and stable
outer energy level (8
electrons), except He
which has 2
This is known as the
octet rule
1
8
2
3
7
Argon
4
6
5
Valence electrons = 8
Ionic Bonds and
Compounds
Writing Chemical Formulas for
Ionic Compounds
Writing Chemical Formulas
for Ionic Compounds
E.Q.: How do you write the
chemical formulas for ionic
compounds?
Ionic Compounds
• Consist of cations (positive ions) and anions
(negative ions)
• Usually composed of metals (cations) and
nonmetals (anions)
• Oxidation numbers (charges of ions) are
important
Types of Ions
• Monatomic ions: Consist of a single
atom (Ex.: Na+, Cl-, Mg+2)
• Mon = Single or one
• Can often be determined from the Periodic
Table
Types of Ions
• Polyatomic ions: Groups of atoms that
behave as a unit and carry a charge
(Ex.: NO3-, OH-, SO4-2)
• Poly = Many or several
• You will need a list of polyatomic ions to
determine the name and formula
nitrate
NO3 -1
chromate
CrO4 -2
chlorate
ClO3 -1
dichromate
Cr2O7 -2
sulfate
SO4 -2
phosphate
PO4 -3
carbonate
CO3 -2
acetate
C2H3O2 -1
hydroxide
OH -1
cyanide
CN -1
ammonium
NH4 +1
Binary Ionic compounds
Writing Formulas
• Made up of two monatomic ions (metal
and nonmetal)
• Ex.: Potassium and Chlorine
• Ex.: Calcium and Bromine
Binary ionic compounds
Writing Formulas
• Transition Metals (Groups 3 through 12)
and some metals in Groups 3A and 4A
(except aluminum, cadmium, silver, and
zinc) can have several oxidation
numbers
• Ex.: Iron (III) and Oxygen
• Ex.: Copper (II) and Oxygen
Ionic Compounds with Polyatomic ions
Writing Formulas
• Ions made up of more than one atom
• Act as individual ions
• Rules used for binary compounds still apply
• Use parentheses when more than one
polyatomic ion is needed and use the
appropriate subscripts outside of the
parentheses
nitrate
NO3 -1
chromate
CrO4 -2
chlorate
ClO3 -1
dichromate
Cr2O7 -2
sulfate
SO4 -2
phosphate
PO4 -3
carbonate
CO3 -2
acetate
C2H3O2 -1
hydroxide
OH -1
cyanide
CN -1
ammonium
NH4 +1
Ionic Compounds with Polyatomic ions
Writing Formulas
• Example: Ammonium ion and chloride
ion
*IWB
• Example: Calcium ion and phosphate ion
*IWB
Practice Problems
1. sodium chloride
2. calcium oxide
3. potassium hydroxide
4. magnesium sulfide
5. copper(II) carbonate
6. aluminum oxide
7. iron(III) oxide
8. sodium carbonate
9. aluminum hydroxide
10. ammonium nitrate
11. zinc nitrate
12. magnesium carbonate
Ionic Bonds and
Compounds
Naming Ionic Compounds
Naming Ionic Compounds
E.Q.: How do you write the
chemical name for an ionic
compound?
Naming ionic compounds:
Rules for naming ionic compounds
1. Name cation first and then the anion
CsBr
+
Cs
cation
-
Br
anion
Naming ionic compounds:
Rules for naming ionic compounds
2.
Monatomic cations use the element name
+1
Cs
3.
=
cesium
Monatomic anions use root of the element and end
with –ide
Br
-1
=
bromide
CsBr = cesium bromide
Elements in Binary Compounds
Element
ide Name
Oxygen
oxide
Phosphorous
phosphide
Nitrogen
nitride
Sulfur
sulfide
Naming ionic compounds:
Rules for naming ionic compounds
4.
This applies to transition metals and some metals
in groups 13 (3A) and 14 (4A) (more than one
oxidation number):
Identify the oxidation number for the transition
metal in the name of the compound
Use a Roman numeral in parentheses after name of
cation
Name
Copper (I)
Copper (II)
Iron (II)
Iron (III)
Chromium (II)
Chromium (III)
Lead (II)
Lead (IV)
Oxidation Number
1+
2+
2+
3+
2+
3+
2+
4+
Naming ionic compounds:
Rules for naming ionic compounds
Examples:
Fe
Fe
2+
3+
2-
and
and
O
O
= FeO =
2-
Iron (II) oxide
= Fe2O3 = Iron (III) oxide
Naming ionic compounds:
Rules for naming ionic compounds
5. If compound has a polyatomic ion,
simply use the name of the ion
+1
-1
NaOH = Na
OH
+1
(NH4 ) 2 S = NH4
S
= sodium hydroxide
-2
= ammonium sulfide
Practice Problems
1. MgCl2 Magnesium chloride
2. LiOH Lithium hydroxide
3. ZnCO3 Zinc carbonate
4. K2S Potassium sulfide
5. FePO4 Iron (III) phosphate
6. Ag3N Silver nitride
7. Mn(CN)2 Manganese (II) cyanide
8. AgC2H3O2 Silver acetate
9. BaI2 Barium iodide
10. PbS2 Lead (IV) sulfide
COVALENT BONDING
E.Q.: WHAT ARE COVALENT BONDS AND
HOW DO ELEMENTS FORM THESE
BONDS?
Covalent Bonds
• Covalent Bonds
– Exist between nonmetals bonded together
– Form when atoms of nonmetals share electrons
– There are no ions involved in covalent bonding
since no element is losing or gaining electrons
– Electrons can be shared equally or unequally
– The unequal sharing results in polar molecules
NAMING COVALENT
COMPOUNDS AND
WRITING THEIR FORMULAS
E.Q.: WHAT ARE THE RULES TO
NAMING COVALENT COMPOUNDS AND
HOW ARE THE FORMULAS WRITTEN?
Naming Covalent Compounds
• Names are usually composed of two
words
- First is the name of the first element
in the formula
- Second is the name of the second
element in the formula, but changing
the ending to –Ide
Naming Covalent Compounds
• If there is more than one atom of an
element in the molecule, then we
need to use prefixes to tell us how
many there are…
Number of Atoms
Prefix
1
Mono-
2
di-
3
tri-
4
tetra-
5
penta-
6
hexa-
7
hepta-
8
octa-
9
nona-
10
deca-
Prefix mono• This prefix can be used for any element,
as long as the element is the second
one in the formula…..
Correct name…
HI = Hydrogen monoiodide
Incorrect name…
HI = Monohydrogen monoiodide
Naming Covalent Compounds
• Some common names to some very
important covalent compounds
Common Name
Molecular
Compound Name
water
dihydrogen
monoxide
Ammonia
nitrogen trihydride
N2 O
nitrous oxide
(laughing gas)
dinitrogen monoxide
NO
nitric oxide
nitrogen monoxide
CH 4
methane
carbon tetrahydride
Formula
H2 O
NH3
Naming Covalent Compounds
• Practice……
P2O5 = diphosphorus pentaoxide
CO = carbon monoxide
CF4 = carbon tetrafluoride
IF5 =
Iodine pentafluoride
Writing Formulas for Covalent
Compounds
• Practice……
Antimony tribromide = SbBr3
Hexaboron silicide = B6Si
Dinitrogen trioxide = N2O3
Phosphorus triiodide = PI3
Metallic Bonds
• Metallic bonds exist
between metals
• Occur when two metals,
usually the same metal,
are bonded together
• “sea of electrons”
• “delocalized electrons”
•
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Regents Chemistry
• Electronegativity
How can we tell really tell which
type of bond we have?
• Electronegativity – is the relative ability
of an atom in a molecule to attract
shared electrons to itself
• This tells us what type of bond we have;
– Covalent, polar covalent or ionic
• Electronegativity values are determined
by measuring the polarities of bonds
between various elements to determine
a specific value for each element
Electronegativity
• Electronegativity values for each element
are obtained by using the Periodic Table
• In fact, there is a general trend in
electronegativity we observe in the
Periodic Table
• Electronegativity values increase across
and up the Periodic Table
– See table on pg. 332
Electronegativity
• We take the difference between the
electronegativity values to determine
exactly what type of bond exists, in
essence the polarity of the bond
See table 12.1
Determining Bond Polarity
• If the difference between the
electronegativity values is:
– 0.0 – 0.5: covalent bond (equal sharing)
– 0.6 – 1.6: polar covalent bond (unequal
sharing)
– 1.7 – up: ionic bond (transferring electrons)
Examples
• Use your Reference Tables to determine
the difference in electronegativity values
and the type of bond for each of the
following:
– H-H
– H-Cl
– H-O
– H-S
• H-F
• NaCl
• O2
• KBr
Worksheet
Regents Chemistry
• Intro to valence electrons
Electrons in an atom
• Electrons surround the nucleus of an atom
in specific energy levels or shells
• Each level can hold only a certain amount
of electrons
• It is an atoms ability to the lose, gain or
share electrons from its outer shell that
determine its reactivity
The outer shell
• The outer shell in an atom contains the
valence electrons
• Valence electrons can be lost, gained or
shared to have eight electrons in the outer
shell
• Each group on the table tells the number of
valence electrons
Periodic Table
• Groups 1, 2, 13, 14, 15, 16, 17, 18 have
1,2,3,4,5,6,7,8 valence electrons,
respectively
• We will not consider the transition metals
• See periodic table
Sharing to reach the Octet Rule
• The octet rule states that an atom cannot
have more than 8 electrons in its outer
shell
• Valence electrons are lost, gained or
shared with other atoms to attain 8
electrons in the outer shell
• Eight valence electrons means a filled and
happy shell - like the noble gases
Nonmetals share
• Nonmetals share electrons to reach eight
valence electrons
• Single, double and triple bonds can be
formed by sharing electrons
Metals + non-metals =
lose/gain e• metals and nonmetals interact by losing
and gaining electrons to reach 8 electrons
(filled outer shell)
• The oxidation states on the periodic table
represent this desire to move electrons
• ex: K+ want to lose 1 electron to reach
noble gas configuration of eight electrons
Lewis structures: your
assignment
• The reading and problems focus on
drawing Lewis structures
• Lewis structures are a means to represent
bond formation between atoms
• Covalent bonded compounds have
different Lewis structures than Ionic
bonded compounds
Example of a Lewis Structure
C
H
H
H C H
H
CH4
Covalent bonds
Regents Chemistry
• Lewis Structures
Lewis Structures
 The Lewis Structure is a representation of a
molecule that shows how the valence
electrons are arranged among the atoms in
a molecule
 We used dots around the elemental symbol
to represent the valence electrons
C
Single Lewis Structure - Practice
 Draw four lone electrons first (if necessary)
them pair them up
 Draw Lewis Structures for the following
atoms
Na
Be
Al
Br
Lewis Structures for Ionic
Compounds
 For Lewis Structures of ionic bonds the
atoms are not joined but draw next to each
other
example:
KBr
+
K [
Potassium loses an
electron to achieve
the noble gas configuration
of Argon
Br ]
Bromine gains an
electron to achieve
the noble gas configuration
of Krypton
Lewis Structures – Covalent Bonds
Hydrogen forms stable molecules when it shares
two electrons
 Two electrons fill Hydrogen’s valence shell
 Helium does not form bonds because its valence
shell is already filled; it is a noble gas
 Second row non-metals Carbon through Fluorine
from stable molecules when surrounded by eight
electrons – the Octet Rule

Lewis Structures – Covalent Bonds
 Valence electrons in covalent bonds can
either be bonding pairs, if involved directly
in the bond or lone pairs if not involved in
the bond
Writing Lewis Structures - Rules
Obtain the total sum of the valence electrons from
all of the atoms
 Use one pair of electrons to form a bond between
each pair of bound atoms. For convenience, a line
(instead of a pair of dots) can be used to indicate
each pair of bonding electrons
 Arrange the electrons to satisfy the duet rule for
hydrogen and the octet rule for second row
non metals

Lewis Structures – Covalent Bonds
 Examples
PH
3
H
l
H– P –H
••
Step 1) 8 total valence e- total
Step 2) Draw one pair of
electrons per bond
8-6 = 2 left
Step 3) Arrange the remaining
electrons according to
octet rule
H
H P H
Lewis Structures – Covalent Bond
Practice Examples
HBr
..
H:Br:
··
CF4
Worksheet
Regents Chemistry
– Ionic Lewis Structures
– Multiple bonds in Lewis Structures
– Polyatomic ion Lewis Structures and
Resonance
Lewis Structures for Ionic
Compounds
 For Lewis Structures of ionic bonds the atoms
are not joined but draw next to each other
example:
KBr
+
K [
Br ]
Potassium loses an
electron to achieve
the noble gas configuration
of Argon
Bromine gains an
electron to achieve
the noble gas configuration
of Krypton
Examples of Ionic Lewis
Structures
 Draw Lewis Structures for the following:
NaCl
LiBr
KI
Multiple Bonds and Lewis
Structures…review first
 We have seen how to draw Lewis Structures
for molecules with single bonds
• For example
NH3
1. Sum the total
valence e2. Subtract number
of bonding e3. Place remaining
valence e-
8 total valence e3 bonds x 2e- = 6 bonding
2 e- left over
H N H
H
Multiple Bonds
 Between atoms of the same element
 Example
• Oxygen
O O
O=O
Just
O = O
Also a Lewis Structure
is called a structural model
Example of Multiple Bonds
Nitrogen
N N
N
N
We now meet the octet rule!
Multiple Bonds
 Between atoms of different elements
 CO2
O
C
O
O = C = O
We must use double bonds to meet the octet rule!
Lewis Structures for Polyatomic
Ions and Resonance Structures
 Read pg. 344 (bottom) to 349 and answer
questions a-g in example
12.4 (pg. 347) and a-i in the Self Check
exercise 12.4 (pg. 348)