Transcript Chapter 7 Ionic and Metallic Bonding

Chapter 7
“Ionic Bonding”
Chemistry
Section 7.1 - Ions
 OBJECTIVES:
Determine
the number of
valence electrons in an
atom of a representative
element.
Section 7.1 - Ions
 OBJECTIVES:
Explain
how the octet rule
applies to atoms of metallic
and nonmetallic elements.
Section 7.1 - Ions
 OBJECTIVES:
Describe
form.
how cations
Section 7.1 - Ions
 OBJECTIVES:
Explain
how anions form.
Valence Electrons are…?
 The
electrons responsible for the
chemical properties of atoms, and
are those in the outer energy level.
 Valence electrons - The s and p
electrons in the outer energy level
 the
 Core
highest occupied energy level
electrons – are those in the
energy levels below.
Keeping Track of Electrons

Atoms in the same column...
Have the same outer electron
configuration.
2) Have the same valence electrons.
1)


The number of valence electrons are
easily determined. It is the group
number for a representative element
Group 2: Be, Mg, Ca, etc.

have 2 valence electrons
Electron Dot diagrams are…
A way of showing & keeping
track of valence electrons.
 How to write them?
 Write the symbol - it
represents the nucleus and
inner (core) electrons
 Put one dot for each valence
electron (8 maximum)
 They don’t pair up until they
have to (Hund’s rule)

X
The Electron Dot diagram for
Nitrogen
Nitrogen has 5 valence
electrons to show.
 First we write the symbol.
Then add 1 electron at a
time to each side.
Now they are forced to pair up.
We have now written the electron dot
diagram for Nitrogen.

N
The Octet Rule
In Chapter 6, we learned that noble gases
are unreactive in chemical reactions
 In 1916, Gilbert Lewis used this fact to
explain why atoms form certain kinds of
ions and molecules
 The Octet Rule: in forming compounds,
atoms tend to achieve a noble gas
configuration; 8 in the outer level is stable
 Each noble gas (except He, which has
2) has 8 electrons in the outer level

Formation of Cations
 Metals
lose electrons to attain a noble
gas configuration.
 They make positive ions (cations)
 If we look at the electron
configuration, it makes sense to lose
electrons:
 Na 1s22s22p63s1 1 valence electron
 Na1+ 1s22s22p6 This is a noble
gas configuration with 8 electrons in
the outer level.
Electron Dots For Cations

Metals will have few valence electrons
(usually 3 or less); calcium has only 2
valence electrons
Ca
Electron Dots For Cations
Metals will have few valence electrons
 Metals will lose the valence electrons

Ca
Electron Dots For Cations
Metals will have few valence electrons
 Metals will lose the valence electrons
 Forming positive ions

2+
Ca
This is named the
“calcium ion”.
NO DOTS are now shown for the cation.
Electron Dots For Cations
Let’s
do Scandium, #21
The electron configuration is:
1s22s22p63s23p64s23d1
Thus, it can lose 2e (making it
2+), or lose 3e- (making 3+)
Sc =
2+
Sc
Scandium (II) ion
Sc =
3+
Sc
Scandium (III) ion
Electron Dots For Cations
Let’s
do Silver, element #47
Predicted configuration is:
1s22s22p63s23p64s23d104p65s24d9
Actual configuration is:
1s22s22p63s23p64s23d104p65s14d10
Ag = Ag+
(can’t lose any more, charges
of 3+ or greater are uncommon)
Electron Dots For Cations
Silver
did the best job it
could, but it did not achieve
a true Noble Gas
configuration
Instead, it is called a
“pseudo-noble gas
configuration”
Electron Configurations: Anions
 Nonmetals
gain electrons to attain
noble gas configuration.
 They make negative ions (anions)
 S = 1s22s22p63s23p4 = 6 valence
electrons
 S2- = 1s22s22p63s23p6 = noble gas
configuration.
 Halide ions are ions from chlorine or
other halogens that gain electrons
Electron Dots For Anions
Nonmetals will have many valence
electrons (usually 5 or more)
 They will gain electrons to fill outer shell.

P
3(This is called the “phosphide
ion”, and should show dots)
Stable Electron Configurations
All atoms react to try and achieve a noble
gas configuration.
 Noble gases have 2 s and 6 p electrons.
 8 valence electrons = already stable!
 This is the octet rule (8 in the outer level
is particularly stable).

Ar
Section 7.2 Ionic Bonds and
Ionic Compounds
 OBJECTIVES:
Explain
the electrical
charge of an ionic
compound.
Section 7.2 Ionic Bonds and
Ionic Compounds
 OBJECTIVES:
Describe
three properties
of ionic compounds.
Ionic Bonding
 Anions
and cations are held together
by opposite charges (+ and -)
compounds are called salts.
 Simplest ratio of elements in an ionic
compound is called the formula unit.
 The bond is formed through the
transfer of electrons (lose and gain)
 Electrons are transferred to achieve
noble gas configuration.
 Ionic
Ionic Compounds
1)Also called SALTS
2)Made from: a CATION
with an ANION (or literally
from a metal combining
with a nonmetal)
Ionic Bonding
Na Cl
The metal (sodium) tends to lose its one
electron from the outer level.
The nonmetal (chlorine) needs to gain one
more to fill its outer level, and will accept the
one electron that sodium is going to lose.
Ionic Bonding
+
Na
Cl
-
Note: Remember that NO DOTS
are now shown for the cation!
Ionic Bond
Negative charges are attracted to positive
charges.
 Negative anions are attracted to positive
cations.
 The result is an ionic bond.
 A three-dimensional crystal lattice of
anions and cations is formed.

Preserve Electroneutrality
 When
ions combine, electroneutrality
must be preserved.
 In the formation of magnesium
chloride,
 2 Cl- ions must balance a Mg2+ ion:
 Mg2+ + 2 Cl→
MgCl2
Ionic Bonding
Let’s do an example by combining
calcium and phosphorus:
Ca

P
All the electrons must be accounted for,
and each atom will have a noble gas
configuration (which is stable).
Ionic Bonding
Ca
P
Ionic Bonding
2+
Ca
P
Ionic Bonding
2+
Ca
Ca
P
Ionic Bonding
2+
Ca
Ca
P
3-
Ionic Bonding
2+
Ca
P
Ca
P
3-
Ionic Bonding
2+
Ca
P
2+
Ca
P
3-
Ionic Bonding
Ca
2+
Ca
P
2+
Ca
P
3-
Ionic Bonding
Ca
2+
Ca
P
2+
Ca
P
3-
Ionic Bonding
2+
Ca
2+
Ca
2+
Ca
P
P
33-
Ionic Bonding
= Ca3P2
Formula Unit
This is a chemical formula, which
shows the kinds and numbers of atoms in
the smallest representative particle of the
substance.
For an ionic compound, the smallest
representative particle is called a:
Formula Unit
Properties of Ionic Compounds
1. Crystalline solids - a regular repeating
arrangement of ions in the solid: Fig.
7.9, page 197


Ions are strongly bonded together.
Structure is rigid.
2. High melting points
 Coordination number- number of ions
of opposite charge surrounding it
- Page 198
Coordination Numbers:
NaCl
Both the sodium
and chlorine have 6
CsCl
Both the cesium
and chlorine have 8
TiO2
Each titanium has
6, and each oxygen
has 3
Do they Conduct?
Conducting electricity means allowing
charges to move.
 In a solid, the ions are locked in place.
 Ionic solids are insulators.
 When melted, the ions can move around.
3. Melted ionic compounds conduct.



NaCl: must get to about 800 ºC.
Dissolved in water, they also conduct (free
to move in aqueous solutions)
- Page 198
The ions are free to move when they are
molten (or in aqueous solution), and thus
they are able to conduct the electric current.
Section 9.1
Naming Ions
OBJECTIVES:
Model
the valence
electrons of metal atoms.
Section 9.1
Naming Ions
 OBJECTIVES:
Identify
the charges on
monatomic ions by using the
periodic table, and name the
ions.
Section 9.1
Naming Ions
 OBJECTIVES:
 Define
a polyatomic ion and write
the names and formulas of the
most common polyatomic ions.
Section 9.1
Naming Ions
 OBJECTIVES:
Identify
the two common
endings for the names of
most polyatomic ions.
Atoms and Ions
 Atoms
are electrically neutral.
 Because
there is the same number of
protons (+) and electrons (-).
 Ions
are atoms, or groups of atoms,
with a charge (positive or negative)
have different numbers of protons
and electrons.
 Only electrons can move, and ions are
made by gaining or losing electrons.
 They
An Anion is…
A
negative ion.
 Has gained electrons.
 Nonmetals can gain electrons.

Charge is written as a superscript on the
right.
1F
Has gained one electron (-ide
is new ending = fluoride)
2O
Gained two electrons (oxide)
A Cation is…
 A positive
ion.
 Formed by losing electrons.
 More protons than electrons.
 Metals can lose electrons
+
K
2+
Ca
Has lost one electron (no
name change for positive ions)
Has lost two electrons
Predicting Ionic Charges
Group 1A: Lose 1 electron to form 1+ ions
H+
Li+
Na+
K+
Rb+
Predicting Ionic Charges
Group 2:
Loses 2 electrons to form 2+ ions
Be2+ Mg2+ Ca2+ Sr2+ Ba2+
Predicting Ionic Charges
B3+
Al3+
Ga3+
Group 13: Loses 3
electrons to form
3+ ions
Predicting Ionic Charges
Neither! Group 14
elements rarely form
ions (they tend to share)
Group 14: Do they
lose 4 electrons or
gain 4 electrons?
Predicting Ionic Charges
N3-
Nitride
P3-
Phosphide
As3- Arsenide
Group 15: Gains 3
electrons to form
3- ions
Predicting Ionic Charges
O2-
Oxide
S2-
Sulfide
Se2- Selenide
Group 16: Gains 2
electrons to form
2- ions
Predicting Ionic Charges
F- Fluoride
Cl- Chloride
Group 17: Gains
Br- Bromide 1 electron to form
1- ions
I- Iodide
Predicting Ionic Charges
Group 18: Stable
noble gases do not
form ions!
Predicting Ionic Charges
Many transition elements
have more than one possible oxidation state.
Note the use of Roman
Iron (II) = Fe2+
numerals to show charges
Iron (III) = Fe3+
Naming cations

Two methods can clarify when more
than one charge is possible:
Stock system – uses roman numerals in
parenthesis to indicate the numerical value
2) Classical method – uses root word with
suffixes (-ous, -ic)
1)

Does not give true value
Naming cations
We will use the Stock system.
 Cation - if the charge is always the same
(like in the main group of metals) just
write the name of the metal.
 Transition metals can have more than one
type of charge.

their charge as a roman numeral
in parenthesis after the name of the metal
(Table 9.2, p.255)
 Indicate
Predicting Ionic Charges
Some of the post-transition elements also
have more than one possible oxidation state.
Tin (II) = Sn2+
Lead (II) = Pb2+
Tin (IV) = Sn4+
Lead (IV) = Pb 4+
Predicting Ionic Charges
Some transition elements have only one
possible oxidation state, such as these three:
Silver = Ag+
Zinc = Zn2+
Cadmium = Cd2+
Exceptions:
 Some of the transition metals
have only one ionic charge:
Do not need to use roman
numerals for these:
Silver is always 1+ (Ag+)
Cadmium
and Zinc are always
2+ (Cd2+ and Zn2+)
Practice by naming these:
 Na+
 Ca2+
 Al3+
 Fe3+
 Fe2+
 Pb2+
 Li+
Write symbols for these:
 Potassium
ion
 Magnesium ion
 Copper (II) ion
 Chromium (VI) ion
 Barium ion
 Mercury (II) ion
Naming Anions
Anions
are always the same
charge
Change the monatomic
element ending to – ide
F
a Fluorine atom will
become a Fluoride ion.
Practice by naming these:
Cl
N3Br2O
3+
Ga
Write symbols for these:
Sulfide
ion
Iodide ion
Phosphide ion
Strontium ion
Polyatomic ions are…
Groups of atoms that stay together and
have an overall charge, and one name.
 Usually end in –ate or -ite


Acetate: C2H3O2-

Nitrate:

Nitrite:
NO3
NO2-
MnO4

Permanganate:

Hydroxide: OH- and Cyanide: CN-?
Know Table 9.3 on page 257
 Sulfate: SO42 Sulfite: SO32
Carbonate: CO32-
 Chromate: CrO42
Dichromate:
Cr2O72-
3 Phosphate: PO4
3 Phosphite: PO3

Ammonium: NH4+
(One of the few positive
polyatomic ions)
If the polyatomic ion begins with H, then combine the
word hydrogen with the other polyatomic ion present:
H+ + CO32- →
HCO3hydrogen + carbonate → hydrogen carbonate ion
Section 9.2 Naming and
Writing Formulas for Ionic
Compounds
 OBJECTIVES:
the rules for naming
and writing formulas for
binary ionic compounds.
Apply
Section 9.2 Naming and
Writing Formulas for Ionic
Compounds

OBJECTIVES:
 Apply the rules for naming and writing
formulas for compounds containing
polyatomic ions.
Writing Ionic Compound
Formulas
Example: Barium nitrate (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2+
(
Ba NO3 ) 2
2. Check to see if charges are
balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance subscripts.
Now balanced.
Not balanced!
= Ba(NO3)2
Writing Ionic Compound
Formulas
Example: Ammonium sulfate (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2. Check to see if charges
are balanced.
( NH4+) SO42-
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance the subscripts.
2
Now balanced.
Not balanced!
= (NH4)2SO4
Writing Ionic Compound
Formulas
Example: Iron (III) chloride (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
Fe3+ Cl-
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance the subscripts.
3
Not balanced!
Now balanced.
= FeCl3
Writing Ionic Compound
Formulas
Example: Aluminum sulfide (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2. Check to see if charges
are balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance the subscripts.
3+
Al
2
2S
3
Not balanced!
Now balanced.
= Al2S3
Writing Ionic Compound
Formulas
Example: Magnesium carbonate (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2. Check to see if charges
are balanced.
Mg2+ CO32They are balanced!
= MgCO3
Writing Ionic Compound
Formulas
Example: Zinc hydroxide (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2+
Zn
2. Check to see if charges are
balanced.
3. Balance charges , if necessary,
using subscripts. Use parentheses if
you need more than one of a
polyatomic ion. Use the criss-cross
method to balance the subscripts.
( OH- )2
Not balanced!
Now balanced.
= Zn(OH)2
Writing Ionic Compound
Formulas
Example: Aluminum phosphate (note the 2 word name)
1. Write the formulas for the
cation and anion, including
CHARGES!
2. Check to see if charges are
balanced.
3+
Al
PO4
3-
They ARE balanced!
= AlPO4
Naming Ionic Compounds
 1.
Name the cation first, then anion
 2.
Monoatomic cation = name of the
element
Ca2+ = calcium ion
 3.
ide
Monoatomic anion = root + Cl- = chloride
CaCl2 = calcium chloride
Naming Ionic Compounds
(Metals with multiple oxidation states)
some metals can form more than one
charge (usually the transition metals)
 use a Roman numeral in their name:

PbCl2 – use the anion to find the charge on the
cation (chloride is always 1-)
Pb2+ is the lead (II) cation
PbCl2 = lead (II) chloride
Things to look for:
1) If cations have ( ), the number
in parenthesis is their charge.
2) If anions end in -ide they are
probably off the periodic table
(Monoatomic)
3) If anion ends in -ate or –ite,
then it is polyatomic
Practice by writing the formula
or name as required…
 Iron (II) Phosphate
 Stannous Fluoride
 Potassium Sulfide
 Ammonium Chromate
 MgSO4
 FeCl3
Practice by writing the formula
for the following:
Magnesium
hydroxide
Iron (III) hydroxide
Zinc hydroxide
Hydrates
 Some
compounds contain H2O in their structure. These compounds are called hydrates.
 The H2O can usually be removed if heated.
 A dot separates water: e.g. CuSO4•5H2O is
copper(II) sulfate pentahydrate.
 A greek prefix indicates the # of H2O groups.
Na2SO4•10H2O sodium sulfate decahydrate
nickel(II) sulfate hexahydrate
NiSO4•6H2O
sodium carbonate monohydrate Na2CO3•H2O
BaCl2•2H2O
barium chloride dihydrate
Prefixes
1 mono
2 di
3 tri
4 tetra
5 penta
6 hexa
7 hepta
8 octa
9 nona
10 deca
Hydrates

Examples:
I. Give the name of the following:
1. CuSO4  5H2O
2. MgCl2  6H2O
3. Na2SO4  10H2O
II. Write the formula for:
1. zinc chloride hexahydrate
2. calcium phosphate dihydrate
3. copper (I) chloride pentahydrate
Summary of Naming and
Formula Writing
 For
naming, follow the flowchartFigure 9.20, page 277
 For writing formulas, follow the
flowchart from Figure 9.22, page
278
Helpful to remember...
1. In an ionic compound, the net ionic
charge is zero (criss-cross method)
2. An -ide ending generally indicates a
binary compound
3. An -ite or -ate ending means there is
a polyatomic ion that has oxygen
4. Prefixes generally mean molecular;
they show the number of each atom
Helpful to remember...
5. A Roman numeral after the
name of a cation is the ionic
charge of the cation