1. Chemical Bonding I: Basic Concepts
Chapter 9
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2. Bonding Forces Intermolecular forces are attractive forces between
molecules.
Intramolecular forces hold atoms together in a
molecule.
“Measure” of intermolecular force
boiling point
melting point
DHvap
DHfus
DHsub
Generally, intermolecular forces are much weaker than intramolecular forces.
11.2

3. Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that
participate in chemical bonding.
Group
# of valence e-
e- configuration 1
ns1 2
ns2 13 3
ns2np1 14 4
ns2np2 15 5
ns2np3 16 6
ns2np4 17 7
ns2np5
9.1

4. Lewis Dot Symbols (electron dot notation)
Consists of the symbol of an element and
one dot for each valence electron in an
atom of an element.
Octet Rule
All elements gain or lose electrons so they
have the same electron configuration as the
nearest noble gas; atoms empty or fill (0 or 8)
their outer shell.

5. Lewis Dot Symbols for the Representative Elements &
Noble Gases
9.1

6. The electrostatic force that holds ions together in an ionic compound.
The Ionic Bond
The electrostatic force that holds ions together in an
ionic compound.
Li+ F - Li + F
[He]
1s22s22p6
1s2
[Ne]
1s22s1
1s22s22p5 Li
Li+ + e-
e- + F - F -
Li+ +
Li+
9.2

7. Ionic compounds
form structures
called crystal
lattices. A
crystal lattice is
a 3-dimensional
arrangement of
ions.

8. Formula Unit The simplest ratio of ions represented
in an ionic compound.
The formula unit for
table salt, sodium
chloride is NaCl.

9. Crystal Lattice Lattice energy
A three dimensional geometric arrangement of
particles. Each positive ion is surrounded by
negative ions, and each negative ion is
surrounded by positive ions.
Lattice energy
The energy required to separate 1 mole of the ions
of an ionic compound. The greater the lattice
energy, the stronger the force of attraction.

10. Properties of Ionic Compounds and Lattice Energy
A crystal lattice arrangement, which involves strong attraction between oppositely charged ions, tends almost always to produce certain properties, such as high melting and boiling points and brittleness.

11. Electrostatic (Lattice) Energy
Lattice energy (E) is the energy required to completely separate one mole of a solid ionic compound into gaseous ions.
Coulomb’s Law: the potential energy (E) between two ions is
directly proportional to the product of their charges and inversely
proportional to the distance of separation between them.
Q+ is the charge on the cation
E = k
Q+Q- r
Q- is the charge on the anion
r is the distance between the ions
9.3

12. Lattice energy (E) increases as Q increases and/or
as r decreases.
compound Lattice energy Q values
MgF , −1
MgO , −2
LiF r F−< r Cl−
LiCl

13. 9.3

14. Why should two atoms share electrons?
A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.
Why should two atoms share electrons?
7e-
7e-
8e-
8e- F F + F
Lewis structure of F2
lone pairs F
single covalent bond
single covalent bond F
9.4

15. Lone pair Also called an unshared electron pair. Is a pair
of electrons that is not involved in bonding and
that belongs exclusively to one atom.
It is usually a lone pair on a central atom that
determines geometry and the type of bond
formed (rather than a lone pair on an outer
atom of a molecule).

16. Lewis structure of water
single covalent bonds
2e-
8e-
2e- H + O + H O H or
Double bond – two atoms share two pairs of electrons
8e-
8e-
8e-
double bonds O C or O C
double bonds
Triple bond – two atoms share three pairs of electrons
8e-
triple bond N
8e- or N
triple bond
9.4

17. Lengths of Covalent Bonds
Bond Lengths
Triple bond < Double Bond < Single Bond
9.4

18. 9.4

19. Covalent Network Solids
Composed of atoms interconnected
by a network of covalent bonds
Examples: SiO2 (quartz)
C (diamond)
Properties:
Brittle, non-conductors of heat
and electricity, very hard, high
melting point.

20. Electronegativity A measure of the ability of an atom
to attract electrons in a chemical
bond.
A relative scale of electronegativity
values is used, with the most electro-
negative element, Fluorine, having
a value of 4.0

21. Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms H F
electron rich
region
electron poor
region
e- poor
e- rich F H d+ d-
9.5

22. Electron Affinity - measurable, Cl is highest
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.
Electron Affinity - measurable, Cl is highest
X (g) + e X-(g)
Electronegativity - relative, F is highest
9.5

23. The Electronegativities of Common Elements
9.5

24. Variation of Electronegativity with Atomic Number
9.5

25. Classification of bonds by difference in electronegativity
Bond Type
0 ~ 0.4
Covalent (nonpolar)
 1.67
Ionic
0 < and <1.67
Polar Covalent
Increasing difference in electronegativity
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
9.5

26. 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
9.5

27. Writing Lewis Structures
Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.
Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.
Complete an octet for all atoms except hydrogen
If structure contains too many electrons, form double and triple bonds on central atom as needed.
9.6

28. 5 + (3 x 7) = 26 valence electrons
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center
Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 3 – Draw single bonds between N and F atoms and complete
octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons F N
9.6

29. 4 + (3 x 6) + 2 = 24 valence electrons
Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – C is less electronegative than O, put C in center
Step 2 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e-
4 + (3 x 6) + 2 = 24 valence electrons
Step 3 – Draw single bonds between C and O atoms and complete
octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24 O C
9.6

30. ( ) - - Two possible skeletal structures of formaldehyde (CH2O) H C O
An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.
formal charge on an atom in a Lewis structure =
total number of valence electrons in the free atom -
total number of nonbonding electrons - 1 2
total number of bonding electrons ( )
The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.
9.7

31. ( ) - -1 +1 H C O = 1 2 = 4 - 2 - ½ x 6 = -1 = 6 - 2 - ½ x 6 = +1
C – 4 e-
O – 6 e-
2H – 2x1 e-
12 e-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12 H C O
formal charge on an atom in a Lewis structure = 1 2
total number of bonding electrons ( )
total number of valence electrons in the free atom -
total number of nonbonding electrons
formal charge on C
= 4 -
2 -
½ x 6 = -1
formal charge on O
= 6 -
2 -
½ x 6 = +1
9.7

32. ( ) - H C O = 1 2 = 4 - 0 - ½ x 8 = 0 = 6 - 4 - ½ x 4 = 0 C – 4 e-H C O
C – 4 e-
O – 6 e-
2H – 2x1 e-
12 e-
2 single bonds (2x2) = 4
1 double bond = 4
2 lone pairs (2x2) = 4
Total = 12
formal charge on an atom in a Lewis structure = 1 2
total number of bonding electrons ( )
total number of valence electrons in the free atom -
total number of nonbonding electrons
formal charge on C
= 4 -
0 -
½ x 8 = 0
formal charge on O
= 6 -
4 -
½ x 4 = 0
9.7

33. Formal Charge and Lewis Structures
For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.
Lewis structures with large formal charges are less plausible than those with small formal charges.
Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.
Which is the most likely Lewis structure for CH2O? H C O -1 +1 H C O
9.7

34. What are the resonance structures of the carbonate (CO32-) ion?
A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. O + - O + -
What are the resonance structures of the
carbonate (CO32-) ion? O C - O C - O C -
9.8

35. Exceptions to the Octet Rule
The Incomplete Octet
Be – 2e-
2H – 2x1e-
4e-
BeH2 H Be
B – 3e-
3F – 3x7e-
24e-
3 single bonds (3x2) = 6
9 lone pairs (9x2) = 18
Total = 24 F B
BF3
9.9

36. Exceptions to the Octet Rule
Odd-Electron Molecules
N – 5e-
O – 6e-
11e- NO N O
The Expanded Octet (central atom with principal quantum number n > 2) S F
S – 6e-
6F – 42e-
48e-
6 single bonds (6x2) = 12
18 lone pairs (18x2) = 36
Total = 48
SF6
9.9

37. Single bond < Double bond < Triple bond
The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy.
Bond Energy
H2 (g)
H (g) +
DH0 = kJ
Cl2 (g)
Cl (g) +
DH0 = kJ
HCl (g)
H (g) +
Cl (g)
DH0 = kJ
O2 (g)
O (g) +
DH0 = kJ O
N2 (g)
N (g) +
DH0 = kJ N
Bond Energies
Single bond < Double bond < Triple bond
9.10

38. Average bond energy in polyatomic molecules
H2O (g)
H (g) +
OH (g)
DH0 = 502 kJ
OH (g)
H (g) +
O (g)
DH0 = 427 kJ
Average OH bond energy = 2
= 464 kJ
9.10

39. Metallic Bonds and Properties of Metals
The bonding in metals is explained by the electron sea model, which proposes that the atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons that surrounds metallic cations.
These delocalized electrons are not held by any specific atom and can move easily throughout the solid.
A metallic bond is the attraction between these electrons and a metallic cation.

40. Metallic Bond, A Sea of Electrons

41. Cross Section of a Metallic Crystal
Metallic Bonds
Cross Section of a Metallic Crystal
nucleus &
inner shell e-
mobile “sea”
of e-
11.6

42. Metallic Bonds and Properties of Metals
Metals generally have extremely high boiling points because it is difficult to pull metal atoms completely
away from the group of cations and attracting electrons.

43. Metallic Bonds and Properties of Metals
Metals are also malleable (able to be hammered into sheets) and ductile (able to be
drawn into wire) because of the mobility of the particles.
The delocalized electrons make metals good conductors of electricity.

44. Metal Alloys
A mixture of elements that has metallic properties is called an alloy.
Alloys can be of two basic types.
A substitutional alloy is one in which atoms of the original metal are replaced by other atoms of similar size.
An interstitial alloy is one in which the small holes in a metallic crystal are filled by other smaller atoms.

45. Metals Form Alloys Metals do not combine with metals.
They form alloys which are solutions
of a metal in a metal. Examples are
steel, brass, bronze and pewter.

46. Dipole dipole-dipole force dipole moment
Equal but opposite charges that are separated by
a short distance.
dipole-dipole force
A force of attraction between polar molecules.
dipole moment
A quantitative measure of the polarity of a bond.

47. What is a dipole-dipole force?
If two neutral molecules, each having a permanent
dipole moment, come together such that their
oppositely charged ends align, they will be attracted
to each other. In a liquid or solid these alignments
are favoured over those where like-charged ends of
the molecules are close together and hence repel
each other.

48. Intermolecular Forces
Dipole-Dipole Forces
Attractive forces between polar molecules
Orientation of Polar Molecules in a Solid
11.2

49. Ion –dipole forces
When an ionic substance dissolves in a
polar solvent (that is, a solvent whose
molecules have a permanent dipole moment)
the majority of the solvent molecules orient
themselves with the oppositely charged end
of the solvent molecule near an ion.
Ion - dipole forces are largely responsible for
the dissolution of ionic substances in water.

50. Intermolecular Forces
Ion-Dipole Forces
Attractive forces between an ion and a polar molecule
Ion-Dipole Interaction
11.2

51. 11.2

52. Two types of interactions:
Dispersion Forces
Attractive forces that arise as a result of temporary
dipoles induced in atoms or molecules.
Induced dipole
The separation of positive and negative charges in an
atom or nonpolar molecule is due to the proximity of
an ion or a polar molecule.
Two types of interactions:
1. ion-induced dipole
2. dipole-induced dipole

53. Intermolecular Forces
Dispersion Forces
Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules
ion-induced dipole interaction
dipole-induced dipole interaction
11.2

54. Induced Dipoles Interacting With Each Other
11.2

55. What type(s) of intermolecular forces exist between each of the following molecules?
HBr
HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules.
CH4
CH4 is nonpolar: dispersion forces. S O
SO2
SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.
11.2

56. Intermolecular Forces
Hydrogen Bond
The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A H … B or
A & B are N, O, or F
11.2

57. Hydrogen Bond
11.2