1. Chemical Bonding
Compounds are formed from chemically bound atoms or ions.
Bonding involves only the valence electrons.

2. Chemical Bonding Ionic Compounds Ionic Radii Lattice Energy
Molecular Compounds
Covalent bonds
Bond Order
Bond Strength
Lewis Structures

3. Lewis Symbols
Lewis symbols show the valence electrons as dots arranged around the atomic symbol.
hydrogen:
sodium:
chlorine: H Na Cl
· ·

4. The Octet Rule
Atoms tend to gain, lose, or share electrons until they have eight valence electrons.

5. Ionic and Molecular Compounds
Formation of sodium chloride: Cl
· ·
Na+ [ ] Cl
· · Na +
Formation of hydrogen chloride: Cl
· · Cl
· · H H +
A metal and a nonmetal transfer electrons to form an ionic compound. Two nonmetals share electrons to form a molecular compound.

6. Ionic Compounds
Ionic compounds consist of a lattice of positive and negative ions.
NaCl:

7. Ionic Bonds
An ionic bond is simply the electrostatic attraction between opposite charges.
Ions with charges Q1 and Q2: Q2 Q1 d d Q E 2 1
The potential energy is given by:

8. Ionic Sizes

9. The Lattice Energy
The lattice energy is the enthalpy change required to separate one mole of an ionic compound into its ions:
NaCl(s) ® Na+(g) + Cl-(g) DH = 788 kJ
Lattice energy increases with increasing ionic charges and decreasing ionic sizes.
KCl(s) ® K+(g) + Cl-(g) DH = 701 kJ
MgCl2(s) ® Mg+2(g) + 2 Cl(g) DH = 3795 kJ

10. The Born-Haber Cycle
The Lattice energy can be calculated with Hess’s law and the following steps:
Na+(g) + e- + Cl(g) H
I1(Na)
E(Cl)
Na+(g) + Cl-(g)
Lattice
energy
Na(g) + Cl(g)
DHf°(Cl,g)
Na(g) + ½Cl2(g)
DHf°(Na,g)
Na(s) + ½Cl2(g)
-DHf°(NaCl,s)
NaCl(s)

11. Estimating Lattice Energy
Arrange with increasing lattice energy:
KCl
NaF
MgO
KBr
NaCl d Q E 2 1
701 kJ
910 kJ K+
Cl
3795 kJ d K+
Br
671 kJ
788 kJ d

12. Molecular Compounds The simplest molecule is H2:
Increased electron density draws nuclei together
The pair of shared electrons constitutes a covalent bond.

13. Lewis Structures
Covalent bonding in a molecule is repre-sented by a Lewis structure.
A valid Lewis structure should have an octet for each atom except hydrogen. H +
H2:
® H H
or H H Cl
· · +
® Cl Cl
· ·
Bonding
electrons
Cl2:
or Cl Cl
· ·
Nonbonding electrons

14. Lewis Structures Draw Lewis structures for: or H F HF: H F H O H
· ·
or H F
· ·
HF:
H F
· ·
· ·
· ·
H O H
· ·
or H O H
· ·
H2O:
H N H H
· ·
or H N H H
· ·
NH3:
H C H H
· ·
or H C H H
CH4:

15. Double and Triple Bonds
Atoms can share four electrons to form a double bond or six electrons to form a triple bond. =
O O
· ·
O2:
· ·
N N 
N2:
The number of electron pairs is the
bond order.

16. Electronegativity
Polarity refers to a separation of positive and negative charge. In a nonpolar bond, the bonding electrons are shared equally:
H2,
Cl2:
In a polar bond, electrons are shared unequally because of the difference in Zeff.
HCl:

17. Electronegativity
Electronegativity refers to the ability of an atom in a molecule to attract shared electrons.
The Pauling scale of electro negativity:

18. Bond Polarity H Cl A polar bond can be pictured using partial charges:+ 
H Cl
 = 0.9
2.1
3.0
Electronegativity
Difference
Bond Type
Nonpolar
Polar
2.0 
Ionic

19. Drawing Lewis Structures
Sum the valence electrons from all atoms. Add one for each negative charge and subtract one for each positive charge.
Draw a skeleton structure with atoms attached by single bonds.
Complete the octets of atoms bound to the central atom.
Place extra electrons on the central atom.
If the central atom doesn’t have an octet, try forming multiple bonds.

20. Drawing Lewis Structures
· ·
Cl C Cl O
· ·
· ·
COCl2
24 ve’s
· ·
· ·
· ·
· ·
· ·
· ·
· ·
· ·
HOCl
14 ve’s
H O Cl
· ·
· ·
· · 
· ·
O Cl O O
· ·
· ·
· ·
· ·
ClO3
26 ve’s
· ·
· ·
· ·
· ·
· ·
H C O H H
· ·
CH3OH
14 ve’s
· ·