BONDING
Bonding

As atoms bond with each other, they
decrease their potential energy, thus
creating more stable arrangements of
matter.
Types of Bonding
The force that holds two atoms together
is called a chemical bond.
 There are 3 types of bonding:
 ionic,
 covalent, and
 metallic.

Keeping Track of Electrons
The number of valence electrons are
easily found by looking up the group
number on the periodic table.
 Group 1A - Li, Na, K, etc. - 1 valence
electron
 Group 2A - Be, Mg, Ca, etc. - 2 valence
electrons

Keeping Track of Electrons
Group 3A - B, Al, Ga, etc. - 3 valence
electrons
 Group 4A - C, Si, Ge, etc. - 4 valence
electrons
 Group 5A - N, P, As, etc. - 5 valence
electrons
 Group 6A - O, S, Se, etc. - 6 valence
electrons

Keeping Track of Electrons
Group 7A - F, Cl, Br, etc. - 7 valence
electrons
 Group 8A - He, Ne, Ar, etc. - 8 valence
electrons (except He has 2 valence
electrons)

Electron (Lewis) Dot Diagrams

Write the symbol.

Put one dot for each
valence electron.

Don’t pair electrons up
until you have to.
X
The Electron Dot Diagram
for Nitrogen

5
Nitrogen has __
valence electrons.
N
Electron Configurations
for Cations
Metals lose electrons to attain noble gas
configuration.
 They make positive ions, cations.

Electron Configurations
for Cations

EXAMPLE: Sodium

Na 1s22s22p63s1

The electron that is removed comes from
the highest energy level.

Na+ 1s22s22p6 - noble gas configuration
Electron Dots For Cations

Calcium has 2 valence electrons.
Ca
Electron Dots For Cations

These will come off,
Ca
Electron Dots For Cations

Forming a positive ion.
+2
Ca
Electron Configurations
for Anions
Nonmetals gain electrons to attain noble
gas configuration.
 This means they want an octet of
electrons, 8 electrons.
 They make negative ions, anions.

Electron Configurations
for Anions

EXAMPLE: Sulfur

S 1s22s22p63s23p4

Sulfur has 6 valence electrons and
needs to gain 2 more to have an octet.

S-2 1s22s22p63s23p6 - noble gas
configuration
Electron Dots For Anions
Phosphorous has 5 valence electrons.
 It will gain 3 electrons to fill the outer
shell.

P
-3
P
Stable Electron Configurations
All atoms react to achieve noble gas
configuration.
 Noble gases, except He, have 2 s
electrons and 6 p electrons, totaling
8 valence electrons.
 They obey the octet rule.

Ar
IONIC
BONDING
Ionic Bonding

Anions and cations are involved in ionic
bonding and are held together by
opposite charges, electrostatic
attraction.
Ionic Bonding
The bond is formed through the transfer
of electrons.
 Electrons are transferred to achieve
noble gas configuration.
 Ionic bonds occur between metals and
nonmetals.

Ionic Bonding
Na Cl
Ionic Bonding
+
Na
Cl
-
Ionic Bonding

All the electrons must be accounted for!
Ca
P
Ionic Bonding
Ca
P
Ionic Bonding
2+
Ca
2P
There is still an unpaired electron for
phosphorus, so another calcium is
needed.
Ionic Bonding
2+
Ca
Ca
2P
Ionic Bonding
2+
Ca
1+
Ca
P
3-
Now Ca has an
unpaired electron, so
another P is needed.
Ionic Bonding
3-
2+
Ca
P
1+
Ca
P
Ionic Bonding
3-
2+
Ca
P
2+
Ca
1P
This P has 2 unpaired electrons, so one
more Ca is needed.
Ionic Bonding
Ca
2+
Ca
2+
Ca
3-
P
1P
Ionic Bonding
Ca
2+
Ca
2+
Ca
3-
P
1P
Ionic Bonding
2+
Ca
2+
Ca
2+
Ca
3-
P
P
3-
Ionic Bonding
Ca3P2
Formula Unit
Ionic Compounds
A compound that is composed of ions is
called an ionic compound.
 Note that only the arrangement of
electrons has changed. Nothing about
the atom’s nucleus has changed.

Properties of Ionic Compounds
Ionic compounds have a crystalline
structure, a regular repeating
arrangement of ions in the solid.
 Even though the ions are strongly
bonded to one another, ionic
compounds are brittle.

Crystalline Structure
Ionic Solids are Brittle
+
+
-
+
+
+
+
-
+
+
Ionic Solids are Brittle

Strong repulsion breaks crystal apart.
- + - +
+ - + - + - +
Properties of Ionic Compounds
The structure is rigid.
 They have high melting points because
of strong forces between ions.
 Conduct electricity in the molten and
dissolved states.

Properties of Ionic Compounds

Any compound that conducts electricity
when melted or dissolved in water is an
electrolyte.
Question

How many valence electrons must an
atom have in its outer energy level in
order to be considered stable?
(eight)
Lattice Energy
The energy required to separate one
mole of the ions of an ionic compound is
called lattice energy, which is
expressed as a negative quantity.
 The greater (that is, the more negative)
the lattice energy is, the stronger the
force of attraction between the ions.

Lattice Energy
Lattice energy tends to be greater for
more-highly-charged ions (those atoms
that have more electrons to give or
those atoms that can take more
electrons).
 Lattice energy also tends to be greater
for small ions.

Question

Between the following ionic compounds,
which would be expected to have the
higher (more negative) lattice energy?
LiF or KBr
(LiF would have the higher lattice
energy because Li is smaller than
K and F is smaller than Br.)
Question

Between the following ionic compounds,
which would be expected to have the
higher (more negative) lattice energy?
NaCl or MgS
(MgS would have the higher lattice
energy because Mg can give away
more electrons than Na and sulfur can
take more electrons than Cl.)
Electronegativity Difference

The electronegativity difference for two
elements in an ionic compound is
greater than or equal to 1.7.
STOP
HERE
COVALENT
BONDING
Covalent Compounds
A molecule is an uncharged group of
two or more atoms held together by
covalent bonds.
 Covalent compounds occur between two
nonmetals or a nonmetal and hydrogen.

How Does H2 Form?

The nuclei repel
+
+
How Does H2 Form?
But they are attracted to electrons.
 They share the electrons.

+
+
Covalent Compounds
The attraction of two atoms for a shared
pair of electrons is called a covalent
bond.
 In a covalent bond, atoms share
electrons and neither atom has an ionic
charge.

Covalent Bonds
Covalent bonds occur between 2
nonmetals because nonmetals hold
onto their valence electrons.
 They can’t give away electrons to bond,
yet, they still want noble gas
configuration.

Covalent Bonds
They get it by sharing valence electrons
with each other.
 By sharing, both atoms get to count the
electrons toward noble gas
configuration.

Covalent Bonding

Fluorine has seven valence electrons.
F
Covalent Bonding

A second fluorine atom also has seven
valence electrons.
F
F
Covalent Bonding

The fluorine atoms will share their lone
electrons.
F
F
Covalent Bonding
F

F
The fluorine atoms are getting closer
together in order to share their lone
electrons.
Covalent Bonding
F F
Covalent Bonding
F F
Covalent Bonding
F F
Covalent Bonding

Both end with full orbitals.
F F
Covalent Bonding

The fluorine on the right has 8 valence
electrons!
F F
8 valence
electrons
Covalent Bonding

The fluorine on the left has 8 valence
electrons!
8 valence
electrons
F F
Single Covalent Bond

A single bond is formed from the
sharing of two valence electrons.
Water
H
Each hydrogen has 1 valence
electron.
Each hydrogen wants 1 more.
Water
O
The oxygen has 6 valence
electrons.
The oxygen wants 2 more.
Water
Hydrogen and oxygen share to make
each other happy.
 The first hydrogen is happy, but the
oxygen still wants one more electron.

HO
Water
A second hydrogen attaches.
 Every atom has full energy levels.

HO
H
Electronegativity Difference

The electronegativity difference for two
elements in a covalent compound is
between 0 and 1.7.
Question

Do atoms that share a covalent bond
have an ionic charge?
(No, the atoms share
electrons and neither atom
has a charge.)
Multiple Bonds
Sometimes atoms share more than one
pair of valence electrons.
 A double bond is when atoms share two
pair of electrons, 4 electrons.
 A triple bond is when atoms share three
pair of electrons, 6 electrons.

Multiple Bonds
Triple bonds are stronger and shorter
than double bonds.
 Double bonds are stronger and shorter
than single bonds.

Carbon Dioxide
CO2 - Carbon is central
atom ( I have to tell you)
 Carbon has 4 valence
electrons
 Wants 4 more
 Oxygen has 6 valence
electrons
 Wants 2 more

C
O
Carbon Dioxide

Attaching 1 oxygen leaves the oxygen
1 electron short and the carbon 3
electrons short
CO
Carbon Dioxide

Attaching the second oxygen leaves
both oxygen 1 short and the carbon 2
short
OC O
Carbon Dioxide

The only solution is to share more
O CO
Carbon Dioxide
O CO
Carbon Dioxide
O CO
Carbon Dioxide
O C O
Carbon Dioxide
O C O
Carbon Dioxide
O C O
Carbon Dioxide
Requires two double bonds
 Each atom gets to count all the atoms in
the bond

O C O
Carbon Dioxide
8 valence
electrons
O C O
Carbon Dioxide
8 valence
electrons
O C O
Carbon Dioxide
8 valence
electrons
O C O
METALLIC
BONDING
Metallic Bonds

The bonding in metals is explained by
the electron sea model, which
proposes that the atoms in a metallic
solid contribute their valence electrons
to form a “sea” of electrons that
surrounds metallic cations.
Metallic Bonds
These delocalized electrons are not
held by any specific atom and can move
easily throughout the solid.
• Metal atoms release their valence
electrons into a sea of electrons shared
by all of the metal atoms. The bond that
results from this shared pool of valence
electrons is called a metallic bond.

Metallic Bonds
Metals hold onto their valence electrons
very weakly.
 Think of them as positive ions floating in
a sea of electrons.

Metallic Bond Properties
Electrons are free to move through the
solid.
 Metals conduct electricity.

+
+ + +
+ + + +
+ + + +
Metallic Bond Properties

Metals generally have extremely high
melting points because it is difficult to
pull metal atoms completely away from
the group of cations and attracting
electrons.
Metallic Bond Properties

Metals are malleable (able to be
hammered into sheets).
Metallic Bond Properties

Metals are also ductile (able to be
drawn into wire) because of the mobility
of the particles.
Malleable
+
+ + +
+ + + +
+ + + +
Malleable

Electrons allow atoms to slide by.
+ + + +
+ + + +
+ + + +
Alloys

A mixture of elements that has metallic
properties is called an alloy.