Bonding – Relationships
between Microscopic
Structure and Macroscopic
Properties
Chemical Bonds
Definition:
Any of several forces, especially ionic, metallic
and covalent bonds, by which atoms or ions
are bound in a molecule or crystal.
Chemical Bond and Valence
Electron
The electrons responsible for the chemical
properties of atoms are those in the outer
energy level: VALENCE ELECTRONS.
– Valence electrons - The electrons in the outer
energy level.
– Inner electrons -those in the energy levels
below.
Keeping Track of Electrons
Atoms in the same column
a. Have the same outer electron configuration.
b. Have the same valence electrons.
c. Easily found by looking up the group number
on the periodic table.
d. Group 2A: Be, Mg, Ca, etc.
2 valence electrons
Electron Dot Diagrams: Lewis
Structures
1. A way of keeping track of
valence electrons.
2. How to write them
3. Write the symbol.
4. Put one dot for each valence
electron
5. Don’t pair up until they have to
X
The Electron Dot diagram for
Nitrogen
Nitrogen has 5 valence
electrons.
 First we write the symbol.
Then add 1 electron at a
time to each side.
Until they are forced to pair up.

N
Write the electron dot diagram
for







Na
Mg
C
O
F
Ne
He
Ionic Compounds
- Ionic Bonds
Electron Configurations for
Cations
1. Metals lose electrons to attain noble gas
configuration.
2. They make positive ions.
3. If we look at electron configuration it makes
sense.
 Na 1s22s22p63s1: 1 valence electron
 Na+ 1s22s22p6 : noble gas configuration
Electron Dots For Cations
 Metals will have few valence electrons
Ca
Electron Dots For Cations
 Metals will have few valence electrons
 These will come off
Ca
Electron Dots For Cations
 Metals will have few valence electrons
 These will come off
 Forming positive ions
+2
Ca
Write the electron
configuration diagram and
orbital
notation
for
 Na






Mg
C
O
F
Ne
He
Electron Configurations for
Anions
1. Nonmetals gain electrons to attain noble
gas configuration.
2. They make negative ions.
3. If we look at electron configuration it
makes sense.
 S 1s22s22p63s23p4: 6 valence electrons
 S-2 1s22s22p63s23p6: noble gas
configuration.
Electron Dots For Anions
 Nonmetals will have many valence
.electrons.
 They will gain electrons to fill outer shell.
P
-3
P
Stable Electron Configuration
1.All atoms react to achieve noble gas configuration.
2.Noble gases have 2 s and 6 p electrons.
3. 8 valence electrons .
4. Also called the octet rule.
Ar
Ionic Bonding
A. Anions and cations are held together by
opposite charges.
B. Ionic compounds are called salts.
C. Simplest ratio is called the formula unit.
D.The bond is formed through the transfer of
electrons.
E. Electrons are transferred to achieve noble
gas configuration.
Ionic Bonding
Na Cl
Ionic Bonding: Lewis Structure
+
Na
Cl
-
Ionic Bonding
 All the electrons must be accounted for!
Ca
P
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
-3
-3
Ionic Bonding
Ca3P2
Formula Unit
I. Properties of Ionic Compounds
a. Crystalline structure.
b. A regular repeating arrangement of ions in
the solid.
c. Structure is rigid.
Ionic Compounds
NaCl:
Ionic compounds
consist of
a lattice
of positive
and negative ions.
Lattice: three
dimensional array of
ions
Crystalline Structure
 1. forms a lattice crystal- a 3-d geometric
structure.
a. each negative ion is surrounded by positive
ions.
b. Lattice energy- the energy
required to break one mole of ions
from the ionic bond.
i. the more – the harder it is to break
Ionic solids are brittle
+
+
-
+
+
+
+
-
+
+
Ionic solids are brittle
 Strong Repulsion breaks crystal apart.
- + - +
+ - + - + - +
Conductivity
1.Conducting electricity is allowing charges to
move.
2.In a solid, the ions are locked in place.
3. Ionic solids are insulators.
4. When melted, the ions can move around.
5. Melted ionic compounds conduct.
6. First get them to 800ºC.
7. ELECTROLYTE-Dissolved in water they
conduct.
Salt water vs salt demo
Properties of Ionic Compounds
Ions are strongly bonded- because of strong
forces between ions they have






High melting points.
High boiling points.
High hardness scale.
Very rigid.
Very brittle.
Do not conduct electricity in solids. Only conduct
electricity in melts or aqueous (water) solutions.
 Do not conduct heat well.
Metals
- Metallic Bonds
Metallic Bonds
 How atoms are held together in the solid.
 Metals hold onto there valence electrons
very weakly.
 Think of them as positive ions floating in a
sea of electrons.
Sea of Electrons
 Electrons are free to move through the solid.
 Metals conduct electricity.
+
+ + +
+ + + +
+ + + +
Malleable
 Hammered into shape (bend).
 Ductile - drawn into wires.
+
+ + +
+ + + +
+ + + +
Malleable
 Electrons allow atoms to slide by.
+ + + +
+ + + +
+ + + +
Properties of Metals
Metals are strongly bonded- because of
strong forces between positive ions and the
“sea” of free electrons surrounding them.







High melting points.
High boiling point.
High hardness scale.
Malleable.
Ductile.
Conduct electricity.
Conduct heat well.
Molecular Compounds
- Covalent Bonds
Covalent Bonds
 How atoms are held together in molecular
compounds.
 Nonmetal elements hold onto valence
electrons tightly.
 All want to gain valence electrons to achieve
octet or [He] configuration.
 The lack of valence electron donor (e.g.
metals) results in the sharing of valence
electrons.
Covalent Bonding
H
and
H
H H
Q: Who’s going to gain?
Who’s going to lose?
H H
They share
them!
Covalent Bond,
strong
Covalent Bonding
H H
Weak interactions
between.
Strong interactions
within.
Properties of Molecular
Compounds
Molecules interact weakly due to lack of
strong forces between them.
 Low melting points. Many molecular compounds
are liquids at room temperature.
 Low boiling points. Lots of molecular compounds
are gases at room temperature.
 Soft: very low hardness scale.
 Solids are brittle.
 Do not conduct electricity in solids nor in aqueous
(water) solutions.
 Do not conduct heat well.