1. Chapter 8: Bonding:General Concepts
Cartoon courtesy of NearingZero.net

2. Chemical Bonding Forces that hold groups of atoms together and make them function as a unit.
Bond Energy Energy required to break a bond

3. Bond Polarity and Dipole Moments
Dipolar Molecules
1. Molecules with a somewhat negative end and a somewhat positive end (a dipole moment)
2. Molecules with preferential orientation in an electric field
3. All diatomic molecules with a polar covalent bond are dipolar

4. Bond Polarity and Dipole Moments
Molecules with Polar Bonds but no Dipole Moment
1. Linear, radial or tetrahedral symmetry of charge distribution
a. CO2 - linear
b. CCl4 – tetrahedral
See table 8.2

5. Ionic Bonding
Ionic bond: the electrostatic force that holds ions together in an ionic compound.
Examples of Ionic Compounds (aka Salts):
NaCl
BaCl2

6. Ionic Bonding Electrons are transferred
Electronegativity differences are
generally greater than 1.7
The formation of ionic bonds is
always exothermic!

7. Determination of Ionic Character
Electronegativity difference is not the final determination of ionic character
Compounds are ionic if they conduct electricity in their molten state

8. Properties of Ionic Compounds
Structure:
Crystalline solids
Melting point:
Generally high
Boiling Point:
Electrical Conductivity:
Excellent conductors, molten and aqueous
Solubility in water:
Generally soluble

9. Coulomb’s Law “The energy of interaction between a pair of ions
E is in Joules
r is the distance between the center of the ions
Q1 and Q2 are the charges of the ions
A negative quantity indicates attraction
A positive quantity indicates repulsion

10. Coulomb’s Law Example:
In solid NaCl the distance between the centers of the ions is 2.76 Å (0.276 nm) Calculate the ionic energy per pair of ions:

11. Sodium Chloride Crystal Lattice
Ionic compounds form solid crystals at ordinary temperatures.
Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions.
All salts are ionic compounds and form crystals.

12. Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.
X (g) + e X-(g)
Electronegativity - F is highest

13. Classification of bonds by difference in electronegativity
Bond Type
Covalent
 2
Ionic
0 < and <2
Polar Covalent
Increasing difference in electronegativity
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-

14. Classify the following bonds as ionic, polar covalent, or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2.
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4
Polar Covalent
N – 3.0
N – 3.0
3.0 – 3.0 = 0
Covalent

15. Formation of Ionic compounds
Stable compounds are formed when nonmetallic elements take electrons from metals.
Atoms usually have a noble gas configuration

16. Formation of Ionic compounds
In general:
When a binary ionic compound is formed
the nonmetal has noble gas configuration
The valence orbitals of the representative metal is emptied

17. Predicting formulas of ionic compounds
The term ionic compounds refers to the solid state of the compound
A collection of positive and negative ions arranged to minimize repulsions and maximize attractions

18. Predicting formulas of ionic compounds
Large electronegativity differences between atoms mean electrons will be transferred

19. Predicting formulas of ionic compounds
Hydrogen typically behaves as a nonmetal
The number of electrons transferred depends on how many each atom needs to gain or lose to achieve noble gas notation

20. Predicting formulas of ionic compounds
EXCEPTIONS:
Tin forms Sn2+ and Sn4+
Lead forms Pb2+ and Pb4+
Bismuth forms Bi3+ and Bi5+
Thallium forms Tl+ and Tl3+

21. Energy and Binary Ionic Compounds
Factors that influence stability and structure
Ionic compounds form because together they have lower energy than the original elements

22. Lattice Energy
The energy released when an ionic solid is formed from its ions
LE is negative (exothermic)
Used as a step to calculate energy of formation

23. E is the potential energy
Q+ and Q- is the charge on the cation and anion
r is the distance between the ions
k is a constant based on the compound
Compound
Lattice Energy
(kJ/mol)
Lattice energy increases as Q increases and/or
as r decreases.
Q: +2,-1
Q: +2,-2
MgF2
MgO
2957
3938
LiF
LiCl
1036
853
r F- < r Cl-

24. Calculating Energy of formation Hf
If we know the steps in the process then we can apply Hess’s law
Because energy is a state function
Break the reaction up into steps
Add them up

25. Estimate Hf for Sodium Chloride
Na(s) + ½ Cl2(g)  NaCl(s)
Lattice Energy
-786 kJ/mol
Ionization Energy for Na
495 kJ/mol
Electron Affinity for Cl
-349 kJ/mol
Bond energy of Cl2
239 kJ/mol
Enthalpy of sublimation for Na
109 kJ/mol
Na(s)  Na(g) kJ
Na(g)  Na+(g) + e kJ
½ Cl2(g)  Cl(g) ½(239 kJ)
Cl(g) + e-  Cl-(g) kJ
Na+(g) + Cl-(g)  NaCl(s) kJ
Na(s) + ½ Cl2(g)  NaCl(s) kJ/mol

26. The energy diagram for the formation of MgO and NaCl.
The Lattice energy to combine Mg2+ and O2- is much more negative than the energy needed for the process that produces Mg2+ and O2- ions.