Chapter 7
“Ionic and Metallic
Bonding”
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 highest occupied energy level
 Core electrons – are those in the
energy levels below.
Keeping Track of Electrons



Atoms in the same column...
1) Have the same outer electron
configuration.
2) Have the same valence electrons.
The number of valence electrons are
easily determined. It is the group
number for a representative element
Group 2A: 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:
2
2
6
2
6
2
1
1s 2s 2p 3s 3p 4s 3d
 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 = Ag1+
(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
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 Bonding
Lets 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.
Metallic Bonds are…
 How
metal atoms are held
together in the solid.
 Metals hold on to their valence
electrons very weakly.
 Think of them as positive ions
(cations) floating in a sea of
electrons: Fig. 7.12, p.201
Sea of Electrons
 Electrons
are free to move through
the solid.
 Metals conduct electricity.
+
+ + +
+ + + +
+ + + +
Metals are Malleable
 Hammered into shape (bend).
 Also ductile - drawn into wires.
 Both malleability and ductility
explained in terms of the
mobility of the valence
electrons
- Page 201
Due to the mobility of the
valence electrons, metals have:
1) Ductility and 2) Malleability
Notice
that the
ionic
crystal
breaks
due to ion
repulsion!
Malleable
Force
+
+ + +
+ + + +
+ + + +
Malleable

Mobile electrons allow atoms to slide
by, sort of like ball bearings in oil.
Force
+ + + +
+ + + +
+ + + +
Ionic solids are brittle
Force
+
+
-
+
+
+
+
-
+
+
Ionic solids are brittle

Strong Repulsion breaks a crystal apart,
due to similar ions being next to each
other.
Force
- + - +
+ - + - + - +
Crystalline structure of metal
 If
made of one kind of atom,
metals are among the simplest
crystals; very compact & orderly
 Note Fig. 7.14, p.202 for types:
1. Body-centered cubic:
–every atom (except those on
the surface) has 8 neighbors
–Na, K, Fe, Cr, W
Crystalline structure of metal
2. Face-centered cubic:
–every atom has 12 neighbors
–Cu, Ag, Au, Al, Pb
3. Hexagonal close-packed
–every atom also has 12 neighbors
–different pattern due to hexagonal
–Mg, Zn, Cd
Alloys
 We
use lots of metals every day,
but few are pure metals
 Alloys are mixtures of 2 or more
elements, at least 1 is a metal
 made by melting a mixture of the
ingredients, then cooling
 Brass: an alloy of Cu and Zn
 Bronze: Cu and Sn
Why use alloys?
Properties are often superior to the pure
element
 Sterling silver (92.5% Ag, 7.5% Cu) is
harder and more durable than pure Ag,
but still soft enough to make jewelry and
tableware
 Steels are very important alloys
– corrosion resistant, ductility, hardness,
toughness, cost

More about Alloys…
7.3, p.203 – lists a few alloys
 Types? a) substitutional alloy- the
atoms in the components are about
the same size
 b) interstitial alloy- the atomic sizes
quite different; smaller atoms fit into
the spaces between larger
 “Amalgam”- dental use, contains Hg
 Table
Atoms and Ions - Naming
 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)
– They have different numbers of protons
and electrons.
 Only
electrons can move, and ions
are made by gaining or losing
electrons.
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
1+
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
H1+ Li1+
Na1+
K1+ Rb1+
Predicting Ionic Charges
Group 2A: Loses 2 electrons to form 2+ ions
Be2+ Mg2+ Ca2+ Sr2+ Ba2+
Predicting Ionic Charges
B3+
Al3+
Ga3+
Group 3A: Loses 3
electrons to form
3+ ions
Predicting Ionic Charges
Neither! Group 4A
elements rarely form
(they
tend to
share)
ions
Group 4A: Do they
lose 4 electrons or
gain 4 electrons?
Predicting Ionic Charges
N3-
Nitride
P3-
Phosphide
As3- Arsenide
Group 5A: Gains 3
electrons to form
3- ions
Predicting Ionic Charges
O2-
Oxide
S2-
Sulfide
Se2- Selenide
Group 6A: Gains 2
electrons to form
2- ions
Predicting Ionic Charges
F1- Fluoride
Cl1- Chloride
Group 7A: Gains
Br1- Bromide 1 electron to form
I1- Iodide
1- ions
Predicting Ionic Charges
Group 8A: Stable
noble gases do not
form ions!
Predicting Ionic Charges
Group B elements: 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:
1) Stock system – uses roman
numerals in parenthesis to
indicate the numerical value
2) Classical method – uses root
word with suffixes (-ous, -ic)
• Does not give true value
Naming cations
 We
will use the Stock system.
 Cation - if the charge is always the
same (like in the Group A metals) just
write the name of the metal.
 Transition metals can have more
than one type of charge.
– Indicate their charge as a roman
numeral in parenthesis after the name
of the metal (Table 9.2, p.255)
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
Group B elements: Some transition elements
have only one possible oxidation state, such
as these three:
Silver = Ag1+
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+ (Ag1+)
–Cadmium and Zinc are always
2+ (Cd2+ and Zn2+)
Practice by naming these:
 Na1+
 Ca2+
 Al3+
 Fe3+
 Fe2+
 Pb2+
 Li1+
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
1F
a Fluorine atom will
become a Fluoride ion.
Practice by naming these:
1Cl
3N
Br1O23+
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

1-

Acetate: C2H3O2

Nitrate: NO31-

Nitrite:

Permanganate: MnO41-

Hydroxide: OH1- and Cyanide: CN1-?
NO21-
Know Table 9.3 on page 257
2-
Sulfate: SO4
2 Sulfite: SO3


Carbonate: CO32-
Phosphate: PO433 Phosphite: PO3


Chromate: CrO422 Dichromate: Cr2O7

Ammonium: NH41+
(One of
theion
few
If the polyatomic
begins with H, positive
then
polyatomic
Writing Ionic Compound
Formulas
Example: Barium
2+
1.
Write
the
(
)
Ba
NO
nitrate (note the 2
3 2
formulas
for
word
name)
2. Check
to see Now
= Ba(NO3)2
Not
the cation
if
charges
are
balanc
3.
Balance
balan
and anion,
balanced.
ed.
charges
,
if
ced!
including
necessary,
Writing Ionic Compound
Formulas
Example: Ammonium
+) SO 2(
NH
1. Write the
4 2
4
sulfate (note the 2 word
formulas
for
Now
2.
Check to
name)
=
(NH4)2SO4
Not
the
cation
balanc
see
if
3.
Balance
balan
and
anion,
ed.
charges ,are
if
ced!
including
balanced.
necessary,
Writing Ionic Compound
Formulas
Example: Iron (III)3+ Fe Cl
1. Write the
chloride (note the 2 3
formulas for
Now
balan
word
name)
2. Check
to
= FeCl3
Not
the cation
see
if
3.
Balance
balan
and anion,
charges ,are
if
ced!
including
balanced.
necessary,
Writing Ionic Compound
Formulas
Example: Aluminum
3+
21.
Write
the
Al
S
sulfide (note the 2 word
2
3
formulas
for
name)
2. Check to
Now
bala
=
Al
S
2 3
Not
the cation
see
if
3.
Balance
balan
and anion,
charges ,are
if
ced!
including
balanced.
necessary,
Writing Ionic Compound
Formulas
Example: Magnesium
2+ CO 2Mg
1.
Write
the
3
carbonate
formulas for
2. Check to
They
the cation
= MgCO3
see if
are
and anion,
charges are
balanced
including
(note the 2 word name)
Writing Ionic Compound
Formulas
Example: Zinc
2+
(
)
1.
Write
the
Zn
OH
hydroxide (note the 2 2
formulas
for
word
name)
balan
2. Check
to see Now
= Zn(OH)2
Not
the cation
if charges
3.
Balance are balan
and anion,
balanced., if
charges
ced!
including
necessary,
Writing Ionic Compound
Formulas
Example: Aluminum
3+
31. Write the
Al
PO
4
phosphate
formulas
for
2. Check to seeThey ARE
the
cation
= AlPO4
if charges are balanced!
and
anion,
balanced.
including
(note the 2 word name)
Naming Ionic Compounds
 1.
Name the cation first, then anion
 2.
Monatomic cation = name of the
element
Ca2+ = calcium ion
 3.
Monatomic anion = root + -ide
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