1. CHEMICAL BONDING Notes by OnyangoNgoye For 2Q & 2T 2013 Class
To play the movies and simulations included, view the presentation in Slide Show Mode.

2. Chemical Bonding What is a Chemical Bond? Ionic Bond Covalent Bond
Coordinate Bond
Hydrogen Bond
Metallic Bond
What are the Structures?
Molecular Structures
Giant Atomic Structures
Giant Ionic Structures

3. Forms of Chemical Bonds
There are 2 main bonding types;
Ionic—complete transfer of 1 or more electrons from one atom to another.
Covalent—some valence electrons shared between atoms.

4. Ionic Bonds
Complete electron transfer from an atom of Sodium to an atom Chlorine.
2 Na(s) + Cl2(g) --->
2 Na Cl-
- Ionic compounds. exists primarily between metals at left of periodic table (Groups 1 and 2 and transition metals) and nonmetals at right (Groups 5, 6,&7).

5. Covalent Bonding
The bond arises from the mutual attraction of 2 nuclei for the same electrons. Electron sharing results.
Bond is a balance of attractive and repulsive forces.

6. Ionic Bonding
Describe the role of the electrons in determining ionic bonds.
Illustrate Ionic bonds using diagrams.

7. Electron Distribution in Molecules
Electron distribution is depicted with Lewis electron dot structures
Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS.
G. N. Lewis

8. Bond and Lone Pairs
Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. • •• H Cl
shared or
bond pair
lone pair (LP)
This is called a LEWIS ELECTRON DOT structure.

9. Note that each atom has a single, unpaired electron.
Bond Formation
A bond can result from a “head-to-head” overlap of atomic orbitals on neighboring atoms. Cl H •• • +
Overlap of H (1s) and Cl (2p)
Note that each atom has a single, unpaired electron.

10. Valence Electrons

11. Giant Ionic Structure of Sodium Chloride.

12. Play by the Rules of the Game
2:8:8
The Duplet & Octet rule.
OCTET RULE

13. Building a Dot Structure
Ammonia, NH3
1. Decide on the central atom (the atom with lowest electron affinity); never H.
Hydrogen atoms are always terminal.
Therefore, N is central
2. Count valence electrons
H = 1 and N = 5
Total = (3 x 1) + 5
= 8 electrons; 4 pairs

14. Building a Dot Structure
3. Form a single bond between the central atom and each surrounding atom H N
4. Remaining electrons form LONE PAIRS to complete octet as needed. H •• N
3 BOND PAIRS and 1 LONE PAIR.
Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair.

15. Sulfite ion, SO32- Step 1. Central atom = S
Step 2. Count valence electrons S = 6
3 x O = 3 x 6 = 18
Negative charge = 2
TOTAL = 26 e- or 13 pairs
Step 3. Form bonds
10 pairs of electrons are
now left.

16. Sulfite ion, SO32-
Remaining pairs become lone pairs, first on outside atoms and then on central atom. • •• O S ••
Each atom is surrounded by an octet of electrons.

17. Carbon IV Oxide, CO2 1. Central atom = _______
2. Valence electrons = __ or __ pairs
3. Form bonds.
This leaves 6 pairs.
4. Place lone pairs on outer atoms.

18. Carbon IV Oxioxide, CO2 4. Place lone pairs on outer atoms.
5. So that C has an octet, we shall form DOUBLE BONDS between C and O.
The second bonding pair forms a pi (π) bond.

19. Double and even triple bonds are commonly observed for C, N, P, O, and S
H2CO
SO3
C2F4

20. Sulphur IV Oxide, SO2 O S 1. Central atom = S
2. Valence electrons = 18 or 9 pairs
3. Form double bond so that S has an octet — but note that there are two ways of doing this.
bring in
left pair
OR bring in
right pair • O S ••

21. Sulfur IV Oxide, SO2 This leads to the following structures.
These equivalent structures are called
RESONANCE STRUCTURES. The true electronic structure is a HYBRID of the two.

22. Urea, (NH2)2CO

23. Urea, (NH2)2CO 1. Number of valence electrons = 24 e-
2. Draw sigma bonds.

24. Urea, (NH2)2CO
3. Place remaining electron pairs in the molecule.

25. Urea, (NH2)2CO
4. Complete C atom octet with double bond.

26. Violations of the Octet Rule
Boron Trifluoride
Central atom = _____________
Valence electrons = __________ or electron pairs = __________
Assemble dot structure
The B atom has a share in only 6 pairs of electrons (or 3 pairs). B atom in many molecules is electron deficient.

27. Sulfur Tetrafluoride, SF4
Violations of the Octet Rule
Sulfur Tetrafluoride, SF4
Central atom =
Valence electrons = ___ or ___ pairs.
Form sigma bonds and distribute electron pairs.
5 pairs around the S atom. A common occurrence outside the 2nd period.

28. Violations of the Octet Rule
Odd # of electrons, NO2
Central atom =
Valence electrons = ___ or ___ pairs.
Form sigma bonds and distribute electron pairs. • •• • •• N O N O • • •• •• •• •• O O • • •• ••

29. Formal Atomic Charges Formal charge= Group no. – 1/2 BEs - LPEs
Definition of Formal Charge:
Formal charge=
Group no. – 1/2 BEs - LPEs

30. Carbon Dioxide, CO2 6 - ( 1 / 2 ) 4 = • O C 4 - ( 1 / 2 ) 8 =

31. Boron Trifluoride, BF3 F B•• • B +1 -1
What if we form a B—F double bond to satisfy the B atom octet?

32. # of bonds between a pair of atoms
Bond Order
# of bonds between a pair of atoms
Double bond
Single bond
Acrylonitrile
Triple bond

33. Bond Polarity
HCl is POLAR because it has a positive end and a negative end.
Cl has a greater share in bonding electrons than does H.
Cl has slight negative charge (-d) and H has slight positive charge (+ d)

34. Bond Polarity
Due to the bond polarity, the H—Cl bond energy is GREATER than expected for a “pure” covalent bond.
BOND ENERGY
“pure” bond 339 kJ/mol calc’d
real bond 432 kJ/mol measured
Difference = 92 kJ. This difference is proportional to the difference in ELECTRONEGATIVITY, .

35. Electronegativity, 
 is a measure of the ability of an atom in a molecule to attract electrons to itself.
Concept proposed by
Linus Pauling

36. Linus Pauling,
The only person to receive two unshared Nobel prizes (for Peace and Chemistry).
Chemistry areas: bonding, electronegativity, protein structure

37. Electronegativity Figure 9.9

38. Bond Polarity Which bond is more polar (or DIPOLAR)? O—H O—F 
OH is more polar than OF
and polarity is “reversed.”

39. Molecular Polarity
Molecules—such as HCl and H2O— can be POLAR (or dipolar).
They have a DIPOLE MOMENT. The polar HCl molecule will turn to align with an electric field.
Figure 9.15

40. Compare CO2 and H2O. Which one is polar?
Polar or Nonpolar?
Compare CO2 and H2O. Which one is polar?

41. Carbon Dioxide
CO2 is NOT polar even though the CO bonds are polar.
CO2 is symmetrical.
+1.5
-0.75
Positive C atom is reason CO2 + H2O gives H2CO3

42. B—F bonds in BF3 are polar. But molecule is symmetrical and NOT polar
Molecular Polarity, BF3
B atom is positive and F atoms are negative.
B—F bonds in BF3 are polar.
But molecule is symmetrical and NOT polar

43. Molecular Polarity, HBF2
B atom is positive but H & F atoms are negative.
B—F and B—H bonds in HBF2 are polar. But molecule is NOT symmetrical and is polar.

44. Is Methane, CH4, Polar?
Methane is symmetrical and is NOT polar.

45. Is CH3F Polar?
C—F bond is very polar. Molecule is not symmetrical and so is polar.

46. Substituted Ethylene C—F bonds are MUCH more polar than C—H bonds.
Because both C—F bonds are on same side of molecule, molecule is POLAR.

47. Substituted Ethene C—F bonds are MUCH more polar than C—H bonds.
Because both C—F bonds are on opposing ends of molecule, molecule is NOT POLAR.