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Covalent Bonding

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Title: Covalent Bonding


1
Covalent Bonding
8.1 Formation of Covalent Bonds 8.2 Dative
Covalent Bonds 8.3 Bond Enthalpies 8.4 Estimation
of Average Bond Enthalpies using Data from
Energetics
2
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions 8.6 Bond
Enthalpies, Bond Lengths and Covalent
Radii 8.7 Shapes of Covalent Molecules and
Polyatomic Ions 8.8 Multiple Bonds 8.9 Covalent
Crystals
3
Formation of Covalent Bonds
4
8.1 Formation of Covalent Bonds (SB p.213)
A. Electron Sharing in Covalent Bonds
electrostatic
The shared electron pair spends most of the time
between the two nuclei.
Overlapping of atomic orbitals ? covalent bond
formation
5
8.1 Formation of Covalent Bonds (SB p.213)
A hydrogen molecule is achieved by partial
overlapping of 1s orbitals
6
8.1 Formation of Covalent Bonds (SB p.214)
Electron density map for covalent compounds
There is substantial electron density at all
points along the internuclear axis.
Compare electron-density-map for ionic compounds
Thus electrons are shared between the two atoms.
7
8.1 Formation of Covalent Bonds (SB p.214)
B. Covalent Bonds in Elements
  • Hydrogen molecule

Dot and cross diagram
8
8.1 Formation of Covalent Bonds (SB p.215)
  • Chlorine molecule
  • Oxygen molecule

9
8.1 Formation of Covalent Bonds (SB p.215)
  • Nitrogen molecule

10
8.1 Formation of Covalent Bonds (SB p.216)
C. Covalent Bonds in Compounds
11
8.1 Formation of Covalent Bonds (SB p.216)
12
8.1 Formation of Covalent Bonds (SB p.216 217)
D. Octet Rule and its Limitations
In forming chemical bonds, atoms tend to achieve
the stable noble gas electronic configuration
with 8 electrons in the valence shell (except
helium which has 2 electrons in the valence
shell) by gaining, losing or sharing of electrons.
13
8.1 Formation of Covalent Bonds (SB p.217)
1. Boron Trifluoride (BF3)
B small atomic size high I.E.s required to
become a cation.
Why doesnt B form ionic compounds with F?
14
8.1 Formation of Covalent Bonds (SB p.207)
2. Phosphorus Pentachloride (PCl5)
There is low-lying vacant d-orbital in P.
Check Point 8-1
Why Phosphorus can expand its octet to form PCl5?
15
Dative Covalent Bonds
16
8.2 Dative Covalent Bonds (SB p.218)
Dative Covalent Bonds
A dative covalent bond is formed by the
overlapping of an empty orbital of an atom with
an orbital occupied by a lone pair of electrons
of another atom.
Remarks(1) The atom that supplies the shared
pair of electrons is known as the donor
while the other atom involved in the dative
covalent bond is known as the acceptor. (2) Once
formed, a dative covalent bond cannot be
distinguished from a normal covalent bond.
17
8.2 Dative Covalent Bonds (SB p.218 219)
A. NH3BF3 molecule
18
8.2 Dative Covalent Bonds (SB p.219)
B. Ammonium Ion (NH4)
19
8.2 Dative Covalent Bonds (SB p.219 220)
D. Aluminium Chloride Dimer (Al2Cl6)
Al relative small atomic size high I.E.s
required to become a cation of 3 charge.
AlCl3
20
8.2 Dative Covalent Bonds (SB p.219 220)
D. Aluminium Chloride Dimer (Al2Cl6)
(a dimer of AlCl3)
Why doesnt Al form ionic compounds with Cl?
Check Point 8-2
21
Bond Enthalpies
22
8.3 Bond Enthalpies (SB p.220)
Bond Enthalpy
Bond enthalpy is the energy associated with a
chemical bond. When a chemical bond is broken or
formed, a certain amount of energy is absorbed
from or released to the surroundings.
23
8.3 Bond Enthalpies (SB p.220)
  • Example
  • Combustion of methane

24
8.3 Bond Enthalpies (SB p.221)
Standard enthalpy changes of combustion of the
homologous series of alkanes and alkanols
25
8.3 Bond Enthalpies (SB p.221)
Bond Dissoication Enthalpy
Bond dissociation enthalpy is the enthalpy change
when one mole of a particular bond in a
particular environment is broken under standard
conditions.
26
8.3 Bond Enthalpies (SB p.221)
Why do successive B.D.E. of C-H differ?
(Average) bond enthalpy E(C-H)
415.5 kJ mol-1
27
8.3 Bond Enthalpies (SB p.222)
Average Bond Enthalpies
Average bond enthalpy is the average of the bond
dissociation enthalpies required to break a
particular chemical bond.
Let's Think 1
28
Estimation of Average Bond Enthalpies using Data
from Enegetics
29
8.4 Estimation of Average Bond Enthalpies using
Data from Enegetics (SB p.223)
A. Derived from the Enthalpy Change of
Atomization of a Compound
Atomization of a compound means the breaking down
of one mole of the gaseous compound into its
constituent atoms in the gaseous state.
30
8.4 Estimation of Average Bond Enthalpies using
Data from Enegetics (SB p.223)
  • Example
  • Atomization of methane

The atomization of methane involves the breaking
of a four C-H bonds. Assume that all four C-H
bonds are equal in strength. The average bond
enthalpy of C-H bonds ΒΌ x (1 662) kJ mol-1
415.5 kJ mol-1
E(C-H) 415.5 kJ mol-1
31
8.4 Estimation of Average Bond Enthalpies using
Data from Enegetics (SB p.223)
  • Two ways to determine the enthalpy change of
    atomization of methane
  • 1. From successive bond dissociation enthalpies
  • 2. From enthalpy cycle and Hesss law

32
8.4 Estimation of Average Bond Enthalpies using
Data from Enegetics (SB p.224 225)
B. Derived from the Enthalpy Changes of
Atomization of Two Compounds
The enthalpy change of atomization of butane
(C4H10) and pentane (C5H12) are 5165 kJ mol-1
and 6337 kJ mol-1 respectively. Find a values
for the bond enthalpies of C-H and C-C based on
the above data.
33
8.4 Estimation of Average Bond Enthalpies using
Data from Enegetics (SB p.224 225)
B. Derived from the Enthalpy Changes of
Atomization of Two Compounds
For butane, 3 E(C-C) 10 E(C-H) 5 165 kJ
mol-1 .(1) For pentane, 4 E(C-C) 12 E(C-H)
6 337 kJ mol-1 .(2) Solving simultaneous
equations (1) and (2), we obtain the following
bond enthalpy values. E (C-H) 412.25 kJ mol-1
E (C-C) 347.5 kJ mol-1
34
Use of Average Bond Enthalpies to Estimate the
Enthalpy Changes of Reactions
35
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.225)
Reaction between ethene and hydrogen
Sum of bond enthalpies of products
36
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.226)
Enthalpy level diagram for the reaction between
ethene and hydrogen
37
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.226)
Reaction between methane and oxygen
Check Point 8-5
38
Bond Enthalpies, Bond Lengths and Covalent Radii
39
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.227)
A. Bond Enthalpies as an Indication of the
Strength of Covalent Bonds
  • Gives a direct measure of the strength of a
    covalent bond? It is the energy required to
    break the bond
  • Not in proportion to the bond order(The number
    of bonding electrons divided by two)

40
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.228)
B. Bond Lengths
  • The distance between the two bonded nuclei
  • Inversely related to bond strength
  • Not constant
  • Depends on the local environment of that
    particular bond
  • Determined experimentally by electron
    diffraction, X-ray diffraction or spectroscopic
    techniques

41
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.228)
C. Relationship between Bond Lengths and Bond
Enthalpies
Bond Bond length (nm) Bond enthalpy (kJ mol-1)
H-H Cl-Cl Br-Br I-I H-F H-Cl H-Br H-I 0.074 0.199 0.228 0.266 0.092 0.127 0.141 0.161 436 242 193 151 565 431 364 299
Any conclusion for the relationship between bond
length bond enthalpy?
Usually a longer bond length corresponds to a
lower value of bond enthalpy (weaker bond).
42
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.228)
Special Situation for F2
Bond Bond Length /nm Bond Enthalpy / kJ
mol-1 F-F 0.142 158 Cl-Cl 0.199 242 Br-Br 0.2
28 193 I-I 0.266 151
Explain why the bond enthalpy of F-F is smaller
than that of Cl-Cl even though the bond length of
F-F is the shortest among the halogens.
43
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.228)
As the size of fluorine atom is very small, the
repulsion between the non-bonding pairs of
electrons on the fluorine atoms weaken the F-F
bond.
44
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.228)
D. Covalent Radii
  • Half the internuclear distance between two atoms
    in a covalently bonded molecule
  • Generally taken as half of the bond length of
    homoatomic covalent molecules (where identical
    atoms are bonded together)

45
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.228)
The covalent radius of an atom is taken as half
of the bond length of a homoatomic molecule
46
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.229)
The covalent radii (in nm) of some elements
47
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.229)
Predicting bond length of A-B if rA rB are known
48
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.229 230)
Calculated and experimentally determined bond
length
Bond Calculated bond length (nm) Experimentally determined bond length (nm)
C-O C-F C-Cl C-Br C-C H-Cl C-H N-Cl 0.150 0.149 0.176 0.191 0.154 0.136 0.114 0.173 0.143 0.138 0.177 0.193 0.154 0.128 0.109 0.174
49
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.229 230)
Calculated and experimentally determined bond
length
Check Point 8-6
Bond Calculated bond length (nm) Experimentally determined bond length (nm)
C-O C-F C-Cl C-Br C-C H-Cl C-H N-Cl 0.150 0.149 0.176 0.191 0.154 0.136 0.114 0.173 0.143 0.138 0.177 0.193 0.154 0.128 0.109 0.174
Let's Think 2
50
Shapes of Covalent Molecules and Polyatomic Ions
51
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.231)
Shapes of Covalent Molecules and Polyatomic Ions
  • Geometric arrangement of atoms within the
    molecules or ions
  • The non-bonding electrons (i.e. the lone pair
    electrons) are not taken into account

52
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.231)
Valence shell electron pair repulsion theory
Electron pairs in the valence shell of the
central atom of a molecule will stay as far apart
as possible to minimize the electrostatic
repulsion between electron pairs in the valence
shell. The electron pairs are oriented with the
maximum separation in space so as to minimize the
coulombic repulsion of electron clouds.
53
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.231)
A. Molecules and Polyatomic Ions without Lone
Pair Electrons on the Central Atom
  • Examples
  • 1. Beryllium Chloride (BeCl2) Molecule
  • 2. Boron Trifluoride (BF3) Molecule
  • 3. Methane (CH4) Molecule
  • 4. Ammonium Ion (NH4)
  • 5. Phosphorus Pentachloride (PCl5) Molecule
  • 6. Sulphur Hexafluoride (SF6) Molecule

54
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.231)
1. Beryllium Chloride Molecule (BeCl2)
Electronic Diagram
Shape in Diagram
Bond angle angle between 2 bonds
Shape in word
Linear
55
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.232)
2. Boron Trifluoride Molecule (BF3)
Electronic Diagram
Shape in Diagram
Shape in word
Trigonal planar
56
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.232)
3. Methane (CH4) Molecule
Electronic Diagram
Shape in Diagram
Shape in word
Tetrahedral
57
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.232)
4. Ammonium Molecule (NH4)
Electronic Diagram
Shape in Diagram
Shape in word
Tetrahedral
58
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.232)
5. Phosphorus Pentachloride (PCl5) Molecule
Electronic Diagram
Shape in Diagram
Shape in word
Trigonal bipyramidal
59
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.232)
6. Sulphur Hexafluoride (SF6) Molecule
Electronic Diagram
Shape in Diagram
Shape in word
Octahedral
60
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.235)
B. Molecules and Polyatomic Ions withLone Pair
Electrons on the Central Atom
  • The valence shell electron pair repulsion theory
    states
  • Electrostatic repulsion decreases in the
    following order
  • Lone pair lone pair gt Lone pair bond pair gt
    Bond pair bond pair

61
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.235)
B. Molecules and Polyatomic Ions withLone Pair
Electrons on the Central Atom
  • Examples
  • 1. Ammonia (NH3) Molecule
  • 2. Water (H2O) Molecule
  • 3. Amide Ion (NH2)

62
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.235)
1. Ammonia (NH3) Molecule
Electronic Diagram
Shape in Diagram
Shape in word
lp-lp repulsion gt lp-bp repulsion gt bp-bp
repulsion
Trigonal pyramidal
63
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.236)
2. Water (H2O) Molecule
Electronic Diagram
Shape in Diagram
lp-lp repulsion gt lp-bp repulsion gt bp-bp
repulsion
Shape in word
V-shaped
64
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.236)
3. Amide Ion (NH2)
Electronic Diagram
Shape in Diagram
lp-lp repulsion gt lp-bp repulsion gt bp-bp
repulsion
Shape in word
V-shaped
65
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.237 238)
66
Multiple Bonds
67
8.8 Multiple Bonds (SB p.238)
Single Bonds
  • A covalent bond with two shared electrons

Multiple Bonds
  • Some atoms share more than two electrons in a
    bond
  • e.g. double bond, triple bond

68
8.8 Multiple Bonds (SB p.239)
Comparison of bond lengths and bond enthalpies
between single and multiple bonds
Bond Bond order Bond length (nm) Bond enthalpy (kJ mol-1)
C C C C C ? C 1 2 3 0.154 0.134 0.120 348 612 837
N N N N N ? N 1 2 3 0.146 0.120 0.110 163 409 944
C O C O 1 2 0.143 0.122 360 743
69
8.8 Multiple Bonds (SB p.239)
Effect of Multiple Bonding on Shapes of Molecules
  • Predict the shapes of molecules or polyatomic
    ions with multiple bonds
  • Examples
  • 1. Ethene (CH2 CH2) Molecule
  • 2. Ethyne (CH ? CH) Molecule
  • 3. Carbon Dioxide (CO2) Molecule
  • 4. Sulphur Dioxide (SO2) Molecule

70
8.8 Multiple Bonds (SB p.239)
1. Ethene (CH2 CH2) Molecule
Electronic Diagram
Shape in Diagram
Shape in word
Trigonal planar
71
8.8 Multiple Bonds (SB p.239)
2. Ethyne (CH ? CH) Molecule
Electronic Diagram
Shape in Diagram
Shape in word
Linear
72
8.8 Multiple Bonds (SB p.240)
3. Carbon dioxide (CO2) Molecule
Electronic Diagram
Shape in Diagram
Shape in word
Linear
73
8.8 Multiple Bonds (SB p.240)
4. Sulphur dioxide (SO2) Molecule
Electronic Diagram
Shape in Diagram
Shape in word
Check Point 8-8
V-shaped
74
Covalent Crystals
75
8.9 Covalent Crystals (SB p.240)
Covalent Crystals
  • May have simple molecular structures or giant
    covalent structures

76
8.9 Covalent Crystals (SB p.240)
Substances with Simple Molecular Structures
  • Consist of discrete molecules held together by
    weak intermolecular forces
  • Atoms in a molecule are held together by strong
    covalent bonds
  • Examples H2 , O2 , H2O, CO2, I2

77
8.9 Covalent Crystals (SB p.240)
Substances with Giant Covalent Structures
  • Consist of millions of atoms bonded covalently
    together in a structural network
  • No simple molecules present
  • Examples diamond, graphite and quartz
    (silicon(IV) oxide)

78
8.9 Covalent Crystals (SB p.240)
1. Diamond
  • Each C atom is covalently bonded to 4 other C
    atoms to form a three-dimensional network
  • The C C bonding pattern accounts for the high
    m.p., stability and extreme hardness
  • Applications scratch proof cookware, watch
    crystals, ball bearings and razor blade

79
8.9 Covalent Crystals (SB p.241)
A diamond crystal
80
8.9 Covalent Crystals (SB p.241)
The structure of diamond
81
8.9 Covalent Crystals (SB p.241)
2. Graphite
  • Each C atom is covalently bonded to 3 other C
    atoms in the same layer. A network of coplanar
    hexagons is formed
  • Weak van der Waals forces hold the layers
    togetherDelocalized e- free to move within layers
  • Properties soft and slippery (used as pencil
    lead), conductor

82
8.9 Covalent Crystals (SB p.241)
Graphite
83
8.9 Covalent Crystals (SB p.241)
The structure of graphite
84
8.9 Covalent Crystals (SB p.242)
Comparison of the properties of diamond and
graphite
Property Diamond Graphite
Density (g cm-3) Hardness Melting point (?C) Colour Electrical conductivity 3.51 10 (hardest) 3 827 Colourless None 2.27 lt 1 (very soft) 3 652 (sublime) Shiny black High
Why graphite has a high m.p. than that of diamond?
85
8.9 Covalent Crystals (SB p.242)
3. Quartz
  • Each Si atom is bonded tetrahedrally to 4
    neighbouring O atoms
  • Each O atom is bonded to 2 Si atoms, one at the
    centre of each of two adjacent tetrahedral
  • Gives rise to a tetrahedral diamondlike structure
    in quartz

86
8.9 Covalent Crystals (SB p.242)
Quartz
87
8.9 Covalent Crystals (SB p.242)
The structure of quartz
88
The END
89
8.1 Formation of Covalent Bonds (SB p.218)
Check Point 8-1
(a) How many lone pair and bond pair electrons
are present in NH3 and H2O molecules respectively?
Answer
  • Ammonia has one lone pair and three bond pairs of
    electrons.
  • Water has two lone pairs and two bond pairs of
    electrons.

90
8.1 Formation of Covalent Bonds (SB p.218)
Check Point 8-1
(b) Nitrogen can only form one chloride, NCl3,
while phosphorus can form two chlorides, PCl3 and
PCl5. Explain briefly.
Answer
91
8.1 Formation of Covalent Bonds (SB p.218)
Check Point 8-1
92
8.1 Formation of Covalent Bonds (SB p.218)
Check Point 8-1
(b) Phosphorus has low-lying d orbitals which
allow it to expand octet (contain more than eight
outermost shell electrons) whereas nitrogen has
not.
93
8.2 Dative Covalent Bonds (SB p.220)
Check Point 8-2
(a) Draw a dot and cross diagram for the
product formed in the reaction between an ammonia
molecule and a hydrogen chloride molecule.
Answer
94
8.2 Dative Covalent Bonds (SB p.220)
Check Point 8-2
(b) There is a dative covalent bond present in a
HNO3 molecule. Draw a dot and cross diagram of
the molecule.
Answer
95
8.2 Dative Covalent Bonds (SB p.220)
Check Point 8-2
(c) State the major difference between an
ordinary and a dative covalent bond.
Answer
(c) A dative covalent bond is covalent bond in
which the shared pair of electrons is supplied by
only one of the bonded atoms, whereas electrons
in an ordinary covalent bond come from both
bonded atoms.
96
8.3 Bond Enthalpies (SB p.222)
Let's Think 1
Why do two atoms bond together? How does covalent
bond strength compare with ionic bond strength?
Answer
They are of similar strength. For example, the
lattice enthalpy of NaCl is 771 kJ mol1 while
the HH bond enthaly is 436 kJ mol1. It is a
misconception that ionic bond must be stronger
(or weaker) than covalent bond.
97
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) Referring to Table 8-2 on page 222, calculate
the enthalpy change for the following reactions
and state whether the reactions are endothermic
or exothermic. (i) Reaction between nitrogen and
hydrogen. N2(g) 3H2(g) ? 2NH3( g)
Answer
98
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) (i) Sum of average bond enthalpies of
products 6 E(N H) 6 ? (388) kJ
mol1 2 238 kJ mol1 ?H 2 252 (2
328) kJ mol1 76 kJ mol1 ? The reaction
is exothermic.
99
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) (ii) Reaction between hydrogen and
chlorine. H2(g) Cl2(g) ? 2HCl(g)
Answer
(a) (ii) H H Cl Cl ? 2H Cl Sum of
average bond enthalpies of reactants E(H H)
E(Cl Cl) (436 242) kJ mol1 678
kJ mol1
100
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) (ii) Sum of average bond enthalpies of
products 2 E(N Cl) 2 ? (431) kJ
mol1 862 kJ mol1 ?H 678 (862) kJ
mol1 184 kJ mol1 ? The reaction is
exothermic.
101
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) (iii) Complete combustion of hydrogen.
Answer
102
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) (iii) Sum of average bond enthalpies of
products 4 E(O H) 4 ? (463) kJ
mol1 1 852 kJ mol1 ?H 1 368 (1
852) kJ mol1 484 kJ mol1 ? The
reaction is exothermic.
103
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) (iv) Complete combustion of ethanol.
Answer
104
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) (iv) Sum of average bond enthalpies of
products 4 E(C O) 6 E(O H) 4 ?
(743) 6 ? (463) kJ mol1 5 750 kJ
mol1 ?H 4 719 (5 750) kJ mol1
1031 kJ mol1 ? The reaction is exothermic.
105
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) (v) Complete combustion of octane.
Answer
106
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(a) (iv) Sum of average bond enthalpies of
products 32 E(C O) 36 E(O H) 32 ?
(743) 36 ? (463) kJ mol1 40 444 kJ
mol1 ?H 32 104 (40 444) kJ mol1
8 340 kJ mol1 ? The reaction is exothermic.
107
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
(b) Calculate the enthalpy change for the
reaction CH4(g) H2O(g) ? CO(g) 3H2(g) using
the following bond enthalpies. E(C H in CH4)
435 kJ mol1 E(C ? O in CO) 1 078 kJ
mol1 E(H H in H2) 436 kJ mol1 E(H O in
H2O) 464 kJ mol1
Answer
108
8.5 Use of Average Bond Enthalpies to Estimate
the Enthalpy Changes of Reactions (SB p.227)
Check Point 8-5
  • CH4(g) H2O ? CO(g) 3H2(g)
  • Sum of average bond enthalpies of reactants
  • 4 E(C H) 2 E(O H)
  • 4 ? (435) 2 ? (464) kJ mol1
  • 2 668 kJ mol1
  • Sum of average bond enthalpies of products
  • E(C ? O) 3 E(H H)
  • 1 078 3 ? (436) kJ mol1
  • 2 386 kJ mol1
  • ?H 2 668 (2 386) kJ mol1 282 kJ mol1

109
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.230)
Let's Think 2
Why does the covalent radius of a given
element change from one compound to another
compound?
Answer
The size of an atom (its covalent radius) is not
fixed. It is because the size of an atom is
determined by its electron cloud which has a
diffuse shape. In different compounds, the
electron cloud of a given atom may vary slightly
due to the different internal environment (i.e.
the atom that is bonded to).
110
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.230)
Check Point 8-6
(a) Predict the approximate bond lengths of Si
H, P H, S H and H Cl from the following
data (Hint Assume that covalent radii are
additive.)
Bond Bond length (nm)
H H Si Si P P (P4) S S (S4) Cl Cl 0.074 0.235 0.221 0.207 0.199
Answer
111
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.230)
Check Point 8-6
112
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.230)
Check Point 8-6
(b) The bond enthalpies of Si H, P H, S H
and H Cl are given in the following
table Assume the actual bond lengths are
very close to that calculated in (a), describe
the relationship between bond length and bond
enthalpy.
Bond Bond enthalpies (kJ mol1)
Si H P H S H Cl H 318 322 338 431
Answer
113
8.6 Bond Enthalpies, Bond Lengths and Covalent
Radii (SB p.230)
Check Point 8-6
(b) The bond enthalpy of a covalent bond is
related to the length. The larger the bond
length, the weaker the attractive force between
the two bonded atoms and the smaller is the bond
enthalpy.
114
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.237)
Example 8-7A
(a) Explain why a molecule of CCl4 is
tetrahedral, but a molecule of NCl3 is trigonal
pyramidal in shape.
Answer
(a) In a CCl4 molecule, there are four bond pairs
of electrons on the central carbon atom. The bond
pairs have to stay as far away as possible. They
take up the shape of a tetrahedron and thus the
molecule is tetrahedral in shape. The four
electron pairs in a NCl3 molecule take up the
shape of a tetrahedron as well. However, one of
the electron pairs is a lone pair and the other
three are bond pairs. The shape of a NCl3
molecule is thus trigonal pyramidal.
115
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.237)
Example 8-7A
(b) Deduce the shape of a molecule of BCl3.
Answer
(b) A BCl3 molecule has six outermost shell
electrons around the central boron atom, forming
three bond pairs. The shape of the BCl3 molecule
is thus trigonal planar.
116
7.5 Ionic Radii (SB p.208)
Example 8-7A
(c) Draw the structures of molecules of XeF2,
XeF4 and XeF6 where Xe is a noble gas element
with eight electrons in its outermost shell.
Answer
117
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.237 238)
Example 8-7B
The following data refer to the molecules NH3,
H2O and HF.
Molecule Bond length (nm) Bond angle
NH3 0.101 107 ?
H2O 0.096 104.5 ?
HF 0.092
118
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.237 238)
Example 8-7B
(a) Briefly explain the variation in bond length.
Answer
119
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.237 238)
Example 8-7B
(b) Explain why the bond angle of H2O is less
than that of NH3.
Answer
(b) This can be explained by the valence shell
electron pair repulsion theory. The central
oxygen atom in H2O has two lone pairs and two
bond pairs of electrons while the central
nitrogen atom in NH3 has one lone pair and three
bond pairs of electrons. The electrostatic
repulsion between electron pairs decreases in
this order lone pair and lone pair gt lone pair
and bond pair gt bond pair and bond pair Thus,
the bond angle of H2O is less than that of NH3.
120
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.237 238)
Example 8-7B
(c) Match the following bond enthalpies to the
bonds in the above three molecules 562 kJ
moll, 388 kJ moll, 463 kJ moll
Answer
121
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.238)
Check Point 8-7
What are the shapes of a H2S molecule and a H3O
ion? Explain their shapes in terms of the valence
shell electron pair repulsion theory.
Answer
H2S molecule is V-shaped. In H2S molecule, there
are two bond pairs and two lone pairs of
electrons in the outermost shell of the central
sulphur atom. All three types of electrostatic
repulsion (lone pair lone pair, lone pair
bond pair, bond pair bond pair) are present.
The two lone pairs will stay the furthest apart
and the separation between the lone pair and a
bond will be greater that that between the two
bond pairs. Therefore, the H S H bond angle
in the H2S molecule is about 104.5? instead of
109.5? in tetrahedron.
122
8.6 Shapes of Covalent Molecules and Polyatomic
Ions (SB p.238)
Check Point 8-7
H3O ion has a trigonal pyramidal shape. In H3O
ion, the central oxygen atom forms two covalent
bonds with two hydrogen atoms respectively. Also,
one dative covalent bond is formed between the
oxygen atom and the remaining hydrogen ion. We
can regard the central oxygen atom has three bond
pairs and one lone pair of electrons. According
to the valence shell electron pair repulsion
theory, the lone pair will stay further away from
the three bond pairs. The three bond pairs are in
turn compressed closer together. Thus, the H O
H bond angles in the H3O ion are about 107?
instead of 109.5? in tetrahedron.
123
8.8 Multiple Bonds (SB p.240)
Check Point 8-8
(a) Does sulphur obey the octet rule in forming a
SO2 molecule? Explain your answer.
Answer
(a) In the formation of SO2 molecule, sulphur
does not obey the octet rule because sulphur has
10 electrons in its outermost shell.
124
8.8 Multiple Bonds (SB p.240)
Check Point 8-8
(b) Draw a dot and cross diagram of the
hydrogen cyanide molecule (HCN). Describe and
explain the shape of the molecule.
Answer
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