1. CHEMISTRY 161
Chapter 9
Chemical Bonding

2. Periodic Table of the Elements
ns2npx
ns1
ns2
chemical reactivity - valence electrons

3. THE OCTET RULE ns2np6 atoms combine to form compounds
in an attempt to obtain a stable
noble gas electron configuration
ns2np6
Iso electronic

4. A + B → AB 1. ELECTRON FULLY TRANSFERED IONIC BONDING
2 Na(s) + Cl2(g)  2 NaCl(s)
2. ELECTRON SHARING
COVALENT BONDING
2 H2(g) + O2(g)  2 H2O(l)

5. represents one valence electron
LEWIS MODEL OF BONDING
LEWIS
DOT SYMBOL
DOT
represents one valence electron H.
Gilbert Lewis
( )

6. . . . . . . . . . .
with the exception of He, the main group number represents number of ‘dots’
only valence electron are considered

7. IONIC BONDING Na electron transfer Ne core implied in symbol
1s22s22p63s1
Lewis Symbol

8. Cl Na Ne core implied in symbol 1s22s22p63s23p5 1s22s22p63s1
Lewis Symbol

9. IONIC BONDING  Cl Cl Na Na+
the formation of ionic bonds is represented
in terms of Lewis symbols Cl Cl Na 
Na+
1s22s22p63s23p6
1s22s22p6
the loss or gain of electrons(dots) until
both species have reached an octet of electrons

10. represents one orbitalCl Cl
[Ne] 3s23p6
represents one orbital
(Pauli: 2 electrons)

11. ions stack together in regular crystalline structures
electrostatic interaction
ionic solids typically
1. high melting and boiling points
2. brittle
3. form electrolyte solutions if they dissolve in water

12. Li(s) + ½ F2(g) → LiF(s) enthalpy of formation
lattice energy
(up to few 1000 kJmol-1)
Li+(g) + F-(g) → LiF(s)
Hess’s Law
Born-Haber Cycle

13. Li(s) + ½ F2(g) → LiF(s) Li+(g) + F-(g) Li(g) + F(g) Li(s) + ½ F2(g)5
ΔHoR= Σ ΔHoi
i=1
ΔHo4
ΔHo3
ΔHo5
Li(g) + F(g)
ΔHo1
ΔHo2
ΔHoR
Li(s) + ½ F2(g)
LiF(s)

14. Mg(s) + ½ O2(g) → MgO(s) Mg2+(g) + O2-(g) Mg+(g) + O-(g) Mg(g) + O(g)
ΔHo6 7
ΔHo5
ΔHoR= Σ ΔHoi
i=1
Mg+(g) + O-(g)
ΔHo7
ΔHo3
ΔHo4
Mg(g) + O(g)
ΔHo1
ΔHo2
ΔHoR
MgO(s)
Mg(s) + ½ O2(g)

15. THE OCTET RULE COVALENT BONDING sharing electrons (electron pair) F
electronic configuration of F is 1s22s22p5 F F

16. F F + F F
non-bonding, or lone pair of electrons
bonding pair of electrons

17. H2 is the simplest covalent molecule+ H H H H +
the bond length of H2 is the distance where the total energy of the molecule is minimum

19. EXAMPLES H2O NH3 CH4 HX single bonds O2 CO2 C2H4 double bonds C2H2 N2
HCN
triple bonds

20. few 100 kJ/mol

22. IONIC OR COVALENT electronegativity difference between two atoms
involved in the bond

23. energies of the atomic orbital with the unpaired electron
ELECTRONEGATIVITY
is the tendency of an atom in a bond to attract shared electrons to itself
F > O > N, Cl > Br > I, C, S …….. Na, Ba, Ra > K, Rb > Cs, Fr
Electronegativity increases F Cl I O N C S Na K Rb Cs Fr Ra Ba Li Br
energies of the atomic orbital with the unpaired electron

24. ELECTRONEGATIVITY F is the most electronegative
F > O > N, Cl > Br > I, C, S …….. Na, Ba, Ra > K, Rb > Cs, Fr
Electronegativity increases Li Br P I N O Cl F C S Na K Rb Cs Fr Ra Ba Se
F is the most electronegative
H has an electronegativity about the same a P

25. IONIC VERSUS COVALENT BONDS
bonds are neither completely ionic nor covalent
(only in homonuclear molecules)

26. IONIC VERSUS COVALENT BONDS
compounds composed of elements with large difference in
ELECTRONEGATIVITY
significant ionic character in their bonding
B has greater electronegativity A B

27. IONIC VERSUS COVALENT BONDS
B has a greater share A B

28. Fluorine is more electronegative than hydrogen.
HYDROGEN FLUORIDE
Fluorine is more electronegative than hydrogen. + F F H H +

29. Fluorine is more electronegative than hydrogen.
HYDROGEN FLUORIDE
Fluorine is more electronegative than hydrogen. + F F H H + d+ d–
This is a polar covalent bond (dipole moment).
The bond has a partly ionic and partly covalent nature.

30. Microwave Spectroscopy
molecules need a dipole moment

31. Variation of ionic character with electronegativity.

32. LEWIS SYMBOLS
IONIC COMPOUDS
COVALENT COMPOUNDS
ELECTRONEGATIVITY

33. single – double – triple
Lewis considers only valence electrons
H2O O H
bonding pair of electrons
non-bonding, or lone pair of electrons
single – double – triple

34. LEWIS STRUCTURES 1. concept of resonances
2. exceptions to the octet rule

35. 1. RESONANCES O O
NO3-
O: 1s22s22p4
N: 1s22s22p3
plus one extra electron for negative charge

36. - + -

37. experiment shows all three bonds are the same
128 pm N
bond angles 120 0
any one of the structures suggests one is different!

38. modify the description by blending the structures
128 pm
bond angles 120 0
modify the description by blending the structures
blending of structures is called resonance

39. use a double headed arrow between the structures
RESONANCE
use a double headed arrow between the structures N O O N N O
electrons involved are said to be DELOCALIZED over the structure.
blended structure is a
RESONANCE HYBRID

40. We use a double headed arrow between the structures..
RESONANCE
We use a double headed arrow between the structures.. N O O N N O O N

41. CO32-
NO2-

42. 2. Exceptions to the octet rule
1. more than 8 electrons around central atom
2. less than an octet around central atom
3. molecules with unpaired electrons

43. 1. more than 8 electrons around central atom
elements in rows 3 and following can exceed octet rule
SF6 S F F S
participation of d electrons

44. Lewis structure for SF6
1s22s22p5
F has seven
S has six
1s22s22p63s22p4
SF SF4 SF6
PF3 PF5
NF3 NF5
ClO SO42-
I3-

45. 2. less than an octet around central atom
BeH2
AlF3
resonances
BF3
NH3
(dative bond)
Lewis base
Lewis acids

46. 3. molecules with unpaired electrons
FREE RADICALS NO
but not NO-