1. Chemistry B2A
Chapter 12
Chemical Bonding

2. Chemical Bonds Ionic bonds 2. Covalent bonds 3. Metallic bonds
4. Hydrogen bonds
5. Van der Waals forces

3. Chemical Bonds
Ionic bonds
2. Covalent bonds

4. Review Level 1 Maximum 2 electrons in valence level
Hydrogen and Helium
main-group elements
1A – 8A
Other levels
Maximum 8 electrons in valence level

5. Na: 1s2 2s2 2p6 3s1 Na+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6
Octet & Duet rules
Noble gases
(Stable)
Goal of atoms
Filled valence level
Na+: 1s2 2s2 2p6
+ e-
Na: 1s2 2s2 2p6 3s1
Ne: 1s2 2s2 2p6
Ar: 1s2 2s2 2p6 3s2 3p6

6. Na: 1s2 2s2 2p6 3s1 Na+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6
Octet & Duet rules
Noble gases
(Stable)
Goal of atoms
Filled valence level
Na+: 1s2 2s2 2p6
+ e-
Na: 1s2 2s2 2p6 3s1
Ne: 1s2 2s2 2p6
+ e-
Cl: 1s2 2s2 2p6 3s2 3p5
Cl-: 1s2 2s2 2p6 3s2 3p6
Ar: 1s2 2s2 2p6 3s2 3p6

7. Mg: 1s2 2s2 2p6 3s2 Mg2+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6
Octet & Duet rules
Noble gases
(Stable)
Goal of atoms
Filled valence level
+ 2e-
Mg: 1s2 2s2 2p6 3s2
Mg2+: 1s2 2s2 2p6
Ne: 1s2 2s2 2p6

8. Mg: 1s2 2s2 2p6 3s2 Mg2+: 1s2 2s2 2p6 Ne: 1s2 2s2 2p6 O: 1s2 2s2 2p4
Octet & Duet rules
Noble gases
(Stable)
Goal of atoms
Filled valence level
+ 2e-
Mg: 1s2 2s2 2p6 3s2
Mg2+: 1s2 2s2 2p6
Ne: 1s2 2s2 2p6
+ 2e-
O: 1s2 2s2 2p4
O2-: 1s2 2s2 2p6
Ne: 1s2 2s2 2p6

9. Cation (Y+): Na+ Li+ Ca2+ Al3+
Metals: lose 1, 2 or 3 e-
Cation (Y+)
Ions
Nonmetals: gain 1, 2 or 3 e-
Anion (X-)
Number of protons and neutrons in the nucleus remains unchanged.
Cation (Y+): Na+ Li+ Ca2+ Al3+
Anion (X-): Cl- F- O-2

10. Transition elements 1A 2A 3A 4A 5A 6A 7A 8A

11. Two problems of Octet rule
Octet rule cannot be used for type II cations (most of the transition and
inner transition elements).
Fe2+
Fe3+
Cu1+
Cu2+
2. Ions of period 1 and 2 elements with charges greater than +2 are unstable. C
C4+
C4- B
B3+
unstable
unstable

12. Ionic bonds
Metal-Nonmetal
Na: 1s2 2s2 2p6 3s1
Cl: 1s2 2s2 2p6 3s2 3p5
Anion
Cation
Na+: 1s2 2s2 2p6
Cl-: 1s2 2s2 2p6 3s2 3p6

13. Sodium (Na)
NaCl
Chlorine (Cl)

14. matter are neutral (uncharged):
total number of positive charges = total number of negative charges
Na+
Cl-
NaCl
Ca2+
Cl-
CaCl2
Al3+
S2-
Al2S3
Ba2+
O2-
Ba2O2
BaO

15. Naming Binary Ionic compounds
name of metal (cation) + name of anion
NaCl sodium chloride
CaO calcium oxide
Cu2O copper(I) oxide cuprous oxide
CuO copper(II) oxide cupric oxide
Naming Polyatomic Ionic compounds
BaCO3 barium carbonate
Li2SO4 lithium sulfate Li2SO3 lithium sulfite

16. Covalent bonds
Nonmetal-Nonmetal
Metalloid-Nonmetal
Sharing of
valence electrons

17. Lewis Dot Structure H He Li Al C N Cl H Or Cl
Lewis Structure
Cl: 1s2 2s2 2p6 3s2 3p5
H: 1s1
He: 1s2
Ar: 1s2 2s2 2p6 3s2 3p6

18. Unshared pair of electrons
(Lone pair) Cl H x
Shared pair of electrons
(bonding pair of electrons)
Only valance electrons are involved in bonding (ionic and covalent bonds).

19. Electronegativity Electronegativity Ionization energy
A measure of an atom’s attraction for the electrons
Electronegativity
Ionization energy

20. Nonpolar covalent bond: electrons are shared equally.
Covalent bonds
Nonpolar covalent bond: electrons are shared equally.
Polar covalent bond: electrons are shared unequally. H Cl δ+ δ-
Dipole

21. Electronegativity Difference Between Bonded Atoms
Electronegativity & bonds
Electronegativity Difference Between Bonded Atoms
Type of Bond
Less than 0.5
Nonpolar Covalent
0.5 to 1.9
Polar Covalent
Greater than 1.9
Ionic H H
2.1 – 2.1 = 0 Nonpolar covalent N H
3.0 – 2.1 = 0.9 polar covalent Na F
4.0 – 0.9 = 3.1 Ionic

22. Covalent compounds H H –
CH4
H C H
H – C – H – H H O O O O –
CH2O – –
H C H
H – C – H
H C H
H – C – H
Correct H H –
NH3
H N H
H – N – H

23. H H H H
H C C H
H C C H
C2H4 H H H H
C – C
C = C H H H H
Correct
H C C H
H C C H
C2H2
H – C – C – H
H – C C – H

24. Resonance When a molecule has more than one Lewis structure. _
CH3COO-
Resonance structures
Resonance structures _

25. Some exceptions to the Octet ruleB H 6

26. Carbon: normally forms four covalent bonds and has no unshared pairs of electrons.
Hydrogen: forms one covalent bond and no unshared pairs of electrons.
Nitrogen: normally forms three covalent bonds and has one unshared pair of electrons.
Oxygen: normally forms two covalent bonds and has two unshared pairs of electrons.
A Halogen: normally forms one covalent bond and has three unshared pairs of electrons. C H N . . . . .
O = . . . . . Cl . .

27. electronegative element + ide
Naming Binary Covalent compounds
Mono – Di – Tri – Tetra – Penta – Hexa – Hepta – Octa – Nona – Deca
Don’t use “mono” for the 1st element.
Drop the “a” when followed by a vowel.
prefixes and name the less electronegative element + prefixed and the name the more
electronegative element + ide
NO2 nitrogen dioxide
N2O4 dinitrogen tetroxide
CCl4 carbon tetrachloride
S2O3 disulfur trioxide

28. VSEPR: Valence-Shell Electron-Pair Repulsion method
VSEPR Model
VSEPR: Valence-Shell Electron-Pair Repulsion method
Bond angle: angle between two atoms bonded to a central atom.
Regions of electron like to be
as far away as possible from the others.

29. Regions of electron density
Four regions of electron density around an atom:

30. Bond Angles & Geometric Structures
Linear molecules
2 regions
Trigonal planar
molecules
3 regions
Tetrahedral arrangement
4 regions

31. Tetrahedral Electron Pair Geometry (Tetrahedral arrangement)
Methane (CH4)
Ammonia (NH3)
Water (H2O)
Tetrahedral
Trigonal Pyramidal
Bent (V Shaped)

32. Tetrahedral Electron Pair Geometry
Unshared electron paires
H2O
CH4
NH3

33. Centers of δ+ and δ- lie at different places (sides).
Polarity
It has polar bonds.
Centers of δ+ and δ- lie at different places (sides). H δ+ – δ- δ- δ+ δ+ δ+
H – C – H
O = C = O – δ- H δ+
nonpolar molecule δ- O δ- = N C H H δ+ H H δ+ H
polar molecule