1
Module
SCHB021
INORGANIC CHEMISTRY 2
PROF KD MODIBANE
Office: N1010
EXT.: 3783
Email: kwena.modibane@ul.ac.za
LECTURER INFO
MR BN MBOKANE
Office: N3019
EXT.: 4771
Email: njabulo.mbokane@ul.ac.za
CONSULTATION TIME
MON, WED & THUR
11H00 – 16H00
2
Course Content
Module Content Topics
1. Foundations:
Introduction to
inorganic chemistry
1.1 Atomic Properties
1.2 Quantum Chemistry
1.3 Shielding
1.4 Periodic Properties
2. Chemical Bonding,
The Elements and
Their Compounds
2.1. Chemical Bonding
(VBT & MOT)
2.2. Transition Metals
2.3. Coordination chemistry
2.4. Isomers
3. Acid-Base Chemistry and
Electrochemistry
3.1. Acids & Bases
3.2. Oxidation & Reducation
3.3. Latimer diagram
3.4. Frost diagram
3.5. Pourbaix diagram
3
Inorganic Chemistry
INTRODUCTION
What is inorganic chemistry?
4
Introduction
What is inorganic chemistry?
 Organic chemistry is:
 the chemistry of life
 the chemistry of hydrocarbon compounds
 C, H, N, O
 Inorganic chemistry is:
 The chemistry of everything else
 The chemistry of the whole periodic Table
(including carbon)
5
Contrast to organic chemistry?
Double and triple bonds studied in organic
chemistry still a feature.
However, inorganic chemistry makes room
for M-M bonds, which can bond quadruples!!
6
PERIODIC TABLE
Recap
 Periods and Groups
 Elements in the periodic table are arranged in
periods (rows) and groups (columns).
 Atomic number increases as you move across a
row or period.
 Periods: Rows of elements are called periods.
7
Recap
PERIODIC TABLE
Dmitri Mendeleev was the first scientist to
construct a periodic table of the elements.
The difference between Mendeleev’s table and
today’s periodic table is that the Modern
periodic table is organized by increasing
atomic number, not increasing atomic weight.
8
Recap
PERIODIC LAW
Mendeleev even went out on a limb and
predicted the properties of 2 (at the time
undiscovered) elements.
He was very accurate in his predictions, which
led the world to accept his ideas about
periodicity and a logical periodic table.
Mendeleev understood the ‘Periodic Law’ which
states:
When arranged by increasing atomic
number, the chemical elements display a
regular and repeating pattern of chemical
and physical properties.
9
Recap
PERIODIC TABLE
A map of the building block of matter
1
1
IA
1
H
Periodic Table
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
1.00797
2
3
4
5
6
7
3
4
Li
Be
6.939
9.0122
11
Na
12
Mg
22.9898 24.305
19
20
5
B
3
IIIB
21
4
IVB
22
23
24
25
26
27
28
85.47
55
87.62
56
88.905
57
91.22 92.906
72
73
Ba
La
Hf
Ta
Tc
95.94
74
[99]
75
W
Re
132.905 137.34 138.91 178.49 180.948 183.85 186.2
87
88
89
104
105
106
107
Fr
Ra
Ac
Ku
[223]
[226]
[227]
[260]
Ru
Co
10
47.90 50.942 51.996 54.9380 55.847 58.9332 58.71
40
41
42
43
44
45
46
Mo
Fe
9
VIIIB
44.956
39
Cs
Mn
8
40.08
38
Nb
Cr
7
VIIB
39.102
37
Zr
V
6
VIB
Sc
Y
Ti
5
VB
Ca
Sr
7
N
8
O
9
F
He
4.0026
10
Ne
10.811 12.0112 14.0067 15.9994 18.9984 20.179
K
Rb
6
C
18
VIIIA
2
Rh
Ni
Pd
11
IB
29
12
IIB
30
Cu
Zn
63.54
47
65.37
48
Ag
Cd
13
Al
14
Si
15
P
16
S
17
Cl
18
Ar
26.9815 28.086 30.9738 32.064 35.453 39.948
31
32
33
34
35
36
Ga
Ge
65.37
49
72.59 74.9216 78.96
50
51
52
In
Sn
As
Sb
Se
Te
Br
Kr
79.909
53
83.80
54
I
Xe
101.07 102.905 106.4 107.870 112.40 114.82 118.69 121.75 127.60 126.904 131.30
76
77
78
79
80
81
82
83
84
85
86
Os
Ir
190.2
108
192.2
109
Pt
Au
Hg
Tl
Pb
Bi
195.09 196.967 200.59 204.37 207.19 208.980
Po
At
Rn
[210]
[210]
[222]
10
Recap
PERIODIC TABLE
 The period number of an element signifies the
highest, unexcited energy level for an electron
in that element.
 Groups
 Elements within a group have similar
chemical and physical properties.
 They have the same outer (valence)
electron arrangement.
11
Recap
PERIODIC TABLE
 Compress view of the periodic table:
Places the Lanthanides and Actinides at the bottom
of the table.
The Periodic Table can be arrange by subshells:
The s-block is Group IA and & IIA, the p-block is
Group IIIA - VIIIA. The d-block is the transition
metals, and the f-block are the Lanthanides and
Actinide metals
12
Recap
PERIODIC TABLE
1
IA
1
18
VIIIA
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
2
3
4
5
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
Metals
10
11
IB
12
IIB
Nonmetals
6
7
13
Recap
 Down the Periodic Table
 Family: Arranged vertically down the periodic table (columns or group,
1- 18 or 1-8 A,B)
 These elements have the same number of electrons in the outer most
shells, the valence shell.
1
IA
1
18
VIIIA
Alkali Family:
1 e- in the valence shell
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
17
VIIA
2
3
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
4
5
Halogen Family:
7 e- in the valence shell
6
7
14
Recap
 Across the Periodic Table
 Periods: Are arranged horizontally across the periodic table (rows 1-7)
 These elements have the same number of valence shells.
1
IA
1
18
VIIIA
2
IIA
13
IIIA
15
VA
16
VIA
17
VIIA
2nd Period
2
3
14
IVA
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
4
5
6th Period
6
7
15
Recap
ATOMIC PROPERTIES
Atomic mass:
Count the number of neutrons + protons
for the atomic mass.
We write the atomic mass as a superscript
in front of the atomic symbol.
The most common isotope of carbon has a
mass of 12, 12C.
16
Recap
ATOMIC PROPERTIES
Atomic number:
Count the number of protons in the
nucleus for the atomic number.
The number of electrons in a neutral atom
is equal to the number of protons in the
nucleus, or the atomic number.
The atomic number of carbon is 6.
17
Recap
ATOMIC PROPERTIES
Isotopes
These are atoms of the same element that
differ in the number of neutrons. Different
isotopes have different natural abundance.
For hydrogen, the isotope 1H has the highest
natural abundance.
e.g., 1H, 2H, 3H
12C, 13C, 14C
18
Recap
ELECTRONS
Each electron in an atom is described by a
unique set of quantum numbers.
n, principal quantum number, values = 1, 2,
3...
l, subsidiary quantum number, values = 0, 1,
2...(n-1)
ml, magnetic quantum number, integral
values between -l and l
ms, spin quantum number, value= -1/2 or 1/2
19
Recap
Atomic orbitals and quantum
mechanics
s, p and d orbitals
20
Recap
Quantum numbers
21
Recap
Example
 Explain, which set of orbitals is defined by
n=4 and l=1? How many orbitals are there in
this set?
 The principal quantum, n= 4, for shell and
subshell, l=1, which correspond to p-orbitals
and the ml= +1, 0, -1 giving three number of
orbitals
Therefore, three 4p orbitals
22
Introduction of Inorganic Chemistry
 Quantum Mechanics (QM)
 Explains & predicts the behavior of matter and
energy at very small scales
 QM plays a significant role in much of today’s
modern technology [computers, cellphones &
Cameras, etc]
Introduction of Inorganic Chemistry
Fundamental Equations of quantum mechanics
Planck
quantization of energy
E = hn
h = Planck’s constant
n = frequency
de Broglie
wave-particle duality
l = h/mv
l = wavelength
h = Planck’s constant
m = mass of particle
v = velocity of particle
Heisenberg
uncertainty principle
Dx Dpx  h/4p
Schrödinger
wave functions

H   E
Dx uncertainty in position
Dpx uncertainty in momentum
H: Hamiltonian operator
: wave function
E : Energy
Introduction of Inorganic Chemistry
Quantum mechanics requires changes in our way of looking at
measurements.
From precise orbits to orbitals:
mathematical functions describing the probable location and
characteristics of electrons
electron density:
probability of finding the electron in a particular portion of
space
Quantization of certain observables occur: Energies can only take
on certain values.
Quantum Numbers and Atomic Wave functions
Each atom of an element, and the electrons thereof,
can mainly be described with respect to a wave
function:
Ψxyz = Rn,l(r) Y(θ, φ)
Angular Functions: They give the shape of the
orbitals and their orientation in space (determined
by l and ml)
Radial Functions: Describe the probability of
finding the electron at a given distance from the
nucleus (determined by n and l).
26
Introduction of Inorganic Chemistry
The radial variation of atomic
orbitals
Nodes: The regions where wave
functions pass through zero
At large distances from the
nucleus, the electron density
(probability of finding an electron)
falls off rapidly.
NB: s orbital has a nonzero
amplitude at the nucleus.
All other orbitals start at
nucleus.
Z
Y
X
27
Introduction of Inorganic Chemistry
Types of Nodes in Orbitals
Z
i) Radial node (aka Spherical
node)
-Ψradial or R(r) = 0
-there are (n-ℓ-1) radial nodes
for each orbital
ii) Angular node (nodal plane)
-Ψangular or Y(θ, φ) = 0
-there are ℓ angular nodes for
each orbital
f q
Y
X
Introduction of Inorganic Chemistry
 Example:
 How many radial and angular nodes are there in
the 4f orbital?
 # of radial nodes = n – l – 1
= 4-3-1
=0
 # of angular nodes = l
=3
Introduction of Inorganic Chemistry
 Typical Question
 Explain radial wave functions of 1s, 2s, and 3s of
hydrogenic orbitals with the aid of a diagram(s).
31
Example
1. Use sketches to explain radial wavefunctions of 1s, 2s,
and 3s of hydrogenic orbitals, and how many radial and
angular nodes are there in the 3s orbital?
(6)
3sradial node = n-l-1
=3-0-1=2
3sangular node = l=0
32
Recap: General Chemistry
Principal quantum number
n = 1, 2, 3, 4 ….
determines the energy of the electron (in a one electron atom) and
indicates (approximately) the orbital’s effective volume
2p 2 me e 4
e2
k
En  

 2
2 2
2rn
nh
n
n=1
2
3
Angular momentum quantum number
l = 0, 1, 2, 3, 4, …, (n-1)
s, p, d, f, g, …..
determines the number of nodal surfaces
(where wave function = 0).
s
Introduction of Inorganic Chemistry
Radial Probability Functions
Introduction of Inorganic Chemistry
Radial distribution function
36
Recap: General Chemistry
If light can behave as particles,why not
particles behave as wave?
Louis de Broglie
The Nobel Prize in Physics 1929
French physicist (1892-1987)
Introduction of Inorganic Chemistry
Louis de Broglie
Particles can behave as wave.
Relation between wavelength l and the
mass and velocity of the particles.
E = hn and also E = mc2,
E is the energy, m is the mass of the
particle, c is the velocity.
Recap: General Chemistry
The Dual Nature of Electrons
Electrons behave as waves and as particles
The energy of a radiation field occurs in
multiples of a basic unit (photon) that depends
on the frequency or wavelength according to:
E = h = hc
λ
Deriving a photon’s momentum from Einstein’s
expression for energy:
p = mc & E = mc2 lead to p = E/c and p = h/λ
Introduction of Inorganic Chemistry
Wave
Particle Duality
E = mc2 = hn
mc2 = hn
p = h /l
{ since n = c/l}
l = h/p = h/mv
This is known as wave particle duality
Introduction of Inorganic Chemistry
De Broglie’s relation
In 1924, Louis de Broglie predicted the relation
p = mv = h
λ
Experimental verification by Davisson & Germer:
Showed that e- beams reflected or scattered from
a crystal gave a diffraction pattern (consisting of a
number of concentric rings).
This phenomenon can be explained only using
the wave nature of electrons.
Introduction of Inorganic Chemistry
Werner Heisenberg
Heisenberg's name will always be associated with
his theory of quantum mechanics, published in
1925, when he was only 23 years.
• It is impossible to specify the exact
position and momentum of a particle,
simultaneously.
• Uncertainty Principle: Dx.Dp  h/4p
where h is Plank’s Constant, a fundamental
constant with the value 6.62610-34 J s.
Introduction of Inorganic Chemistry
The uncertainity principle
Heisenberg’s uncertainity principle
h = ET = pλ
This relation applied to wave and particle
properties, where T is the period of wave, T=
1/v)
This equation, (Δpx) (Δx) ≥ h, also states that
the product of the uncertainty in momentum with
respect to a given coordinate and the
uncertainty in position with respect to the same
coordinate must be equal or greater than
Plank’s constant.
Introduction of Inorganic Chemistry
∆x . ∆p ≥ ћ
2
∆E . ∆t ≥ ћ
2
where,
∆x is uncertainty in position,
∆p is uncertainty in momentum,
∆t is uncertainty in time, and
∆E is the uncertainty in energy
NB: ћ = h, h is Planck’s constant = 6.626 × 10-34 J.s
2π
Introduction of Inorganic Chemistry
Wave equation?
Schrödinger Equation.
• Most significant feature of the Quantum
Mechanics: Limits the energies to
discrete values, ergo Quantization.
1887-1961
Introduction of Inorganic Chemistry
The Schrödinger Equation
Recap: General Chemistry
Assignment of electrons to orbitals
The aufbau principle: Electron configurations
are built up from the bottom, using the lowest
energy orbitals first.
47
Recap: General Chemistry
Assignment of electrons to orbitals
Hund’s rule: Where orbitals are available in
degenerate sets, electrons are not paired until
each orbital in a degenerate set has been halffilled.
Pauli’s exclusion principle: No two electrons
may have the same set of four quantum
numbers.
 Where two electrons occupy the same orbital,
they must have opposite spins:
ms = + 1/2 for one electron and ms = - 1/2 for the 2nd
electron.
48
Introduction of Inorganic Chemistry
Interactions within atoms
As e-s are added to an atom, they interact
with each other as well as with the nucleus
Importantly, two e-s in the same orbital
have a higher energy than when they
occupy different orbitals.
rule of maximum multiplicity.
Multiplicity = n + 1,
where n is the number of unpaired electrons
49
Consequence of Hund’s Rule
When two negatively charged electrons occupy the
same region of space (same orbital) in an atom,
they repel each other with a Coulombic energy of
repulsion , ∏c, per pair of electrons.
NB: This repulsive force favors electrons in different
orbitals over electrons in the same orbitals.
Additionally, there is an exchange energy, ∏e,
which depends on the number of possible
exchanges between two electrons with the same
energy and the same spin.
50
EXAMPLE
The electron configuration of carbon is
1s22s22p2. The 2p electrons can be placed
in the p orbitals in three ways:
State (1) involves ∏c because it is the only
one that pairs electrons in the same orbital.
The energy of this state is higher than that of
the other two by ∏c.
51
 State (3): the electrons have the same spin & are
therefore indistinguishable from each other. So
there are two possible ways in which the electrons
can be arranged:
 The energy involved in such an exchange of
parallel electrons is designated ∏e ; each
exchange stabilizes (lowers the energy of) an
electronic state, favoring states with more parallel
spins (Hund’s rule).
 State (3), which is stabilized by one exchange of
parallel electrons, is lower in energy than state (2)
by ∏e.
52
∏e = K x P,
where K is a
constant and
P= n(n-1)/2.
53
Inorganic Chemistry
Shielding and Penetration
Each electron acts as a shield for electrons farther
from the nucleus, reducing the attraction between the
nucleus and the more distant electrons.
The presence of an electron inside shells of other
electrons is called penetration.
For outer orbitals, the increasing ∆E between
levels with the same n but different ℓ values forces
overlap of energy levels with n = 3 and n = 4.
Therefore,
4s fills before 3d
5s fills before 4d
6s
5d
4f
5d
5f fills before 6d
54
Inorganic Chemistry
Shielding
As Z increases, e-s are “drawn in” towards the
nucleus, causing orbital energies to become
more negative.
Although energies decrease with increasing
Z, the changes are irregular because of
shielding of outer e-s by inner ones.
Therefore, shielding is the reduction of the
true nuclear charge to the effective nuclear
charge by the other e-s.
55
Effective Nuclear Charge (Z* or Zeff)
A measure of the nuclear attraction for an
electron which is calculated using Z* = Z – S
where Z is the nuclear charge (i.e. atomic #)
S is the shielding constant (can also use σ
instead of S).
NB: The closer to the nucleus that an electron
can approach, the closer is Z* to Z itself.
56
Inorganic Chemistry
Effective Nuclear Charge (Z* or Zeff)
(Z*), can also be explained as the amount of
positive charge felt by the atom’s outer
electrons.
(Z*) depends on the values of n and l of
electron.
(Z*) is always less than the full nuclear
charge.
The negative charge of the electrons in inner
shells partially neutralizes the positive charge
of nucleus
57
Inorganic Chemistry
Slater’s rules for calculating the shielding
constant
Write the electronic configuration of an atom
by grouping the shells as follows:
e.g., (1s), (2s,2p) (3s,3p) (3d) (4s,4p) (4d)
(5s,5p) (5d) (6s,6p) etc.
NB: Electrons in higher groups (to the right in
the list above) do not shield those in lower
groups.
58
Inorganic Chemistry
Slater’s rules for calculating the shielding
constant
 For ns or np valence electrons:
 Electrons in the same ns, np group
contribute 0.35
Except the 1s, where 0.30 works better
 Electrons in the n-1 group contribute 0.85
 Electrons in the n-2 or lower groups
contribute 1.00
59
Inorganic Chemistry
Slater’s rules for calculating the shielding
constant
 For nd and nf valence electrons:
 Electrons in the same nd or nf group
contribute 0.35
 Electrons in groups to the left
contribute1.00
60
Inorganic Chemistry
Example:
Calculate the effective nuclear charge, (Z*) felt
by one of the valence electrons of Nitrogen.
N7 = (1s2) (2s22p3)
Electrons in the n-1 group contribute 0.85
Two electrons in 1s: 0.85 x 2
Electrons in the same ns, np group contribute
0.35
Four electrons in ns, np: 0.35 x 4
61
Inorganic Chemistry
Shielding constant, S = (0.85 x 2) + (0.35 x 4)
= 3.10
Z* = Z – S
= Z- σ
= 7 – 3.10
= 3.9
62
Inorganic Chemistry
Application of Slater’s rules (Z*)
Estimate Z* for one of the 4s electrons, and one of the 3d
electrons of Zn.
Solution:
Grouping: (1s2)(2s2,2p6)(3s2,3p6)(3d10)(4s2)
For the 4s e-:
S = (10×1.00) + (18×0.85) + (1×0.35)
= 25.65
Z* = Z – S
= 30 – 25.65
= 4.35
For the 3d e-: S = (18×1.00) + (9×0.35) = 21.15
Z* = 30-21.15 = 8.85
63
Inorganic Chemistry
Application of Slater’s rules (Z*)
Therefore during ionization, a 4s electron is
lost first, since it has less effective nuclear
charge.
 Energy is also related to Z* as follows:
E = -(Z*)2 (13.6 eV)
n2
= -(Z – S)2 (13.6 eV)
n2
64
Inorganic Chemistry
Application of Slater’s rules (Z*)
Predict the energy of the 4s electron in K
E = -(Z*)2 (13.6 eV)
n2
= -(19 – 16.8)2 (13.6 eV)
42
= -(2.2)2 (13.6 eV)
16
= -4.1 eV(Experimental measurement= -4.3 eV)
65
Inorganic Chemistry
Problem
• Use Slater’s rules to explain why the valence electron
configuration of the ground state of a Cu atom is likely to
be 3d104s1 rather than 3d94s2.
(7)
Cu=1s22s22p63s23p63d104s1 
Z*4s1=29-[(0 x 0.35)+(18 x 0.85)+(10 x 1) 
= 3.70 
Z*3d10=29-[(9 x 0.35)+(18 x 1) 
= 7.85 
The 3d electron feels a greater force of attraction by the
nucleus than the 4s electron. It is therefore easy to ionize
the 4s electron than the 3d electron thus making the
valence configuration of Cu atom to be 3d104s1 
66
67
Recap: General Chemistry
Periodic Properties & Z*
 Rationalize the trends in IE, EA, and χ as
influenced by Z* or Zeff
Also do Atomic v/s Ionic radii
NOTE
Cations are smaller than parent atoms
Anions are larger than parent atoms
 Mg2+, F-, and Na+ all have e- config [Ne] or
[He]2s22p6 said to be isoelectronic
Recap: General Chemistry
Periodic Trends
Ionization energy the energy required to remove
an electron from a gaseous atom or ion of an
atom.
Increases from left to right across a period but
decreases from top to bottom in a group, why?
69
Recap: General Chemistry
Electron attachment enthalpies (Electron
affinity)
The enthalpy change DHEA that accompanies
addition of an electron(s) provides a measure
of the willingness of an atom to form anions.
Where these enthalpy changes are negative,
formation of the anion is favourable
(exothermic).
X (g) + e-  X- (g)
Cl (g) + e-  Cl- (g)
Recap: General Chemistry
Periodic properties of Elements
Electron affinity: energy required to remove an
electron from a negative ion.
Increases upward for the groups and from L to R
across periods of a periodic table
Recap: General Chemistry
Periodic properties of Elements
Atomic size/radii: decreases gradually across
each period.
HOMEWORK
• Read on Scandide Contraction, as well
as Lanthanide Contraction and their
effects on atomic radii and other trends!!
73
Recap: General Chemistry
Periodic properties of Elements
Covalent radius (rcov):
Half the internuclear separation between 2
singly-bonded atoms of the same element.
Van der Waals radius (rvdw):
 Half of the internuclear separation of 2
non-bonded atoms of the same element on
their closest possible approach.
Recap: General Chemistry
Periodic properties of Elements
Covalent v/s Van der Waals Radius
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Inorganic Chemistry
Electronegativity 
Electronegativity is the power of an atom of
the element to attract electrons to itself when
it is part of a compound.
Periodic trends in electronegativity can be
related to the size of the atoms and electron
configuration.
If an atom is small and has an almost closed
shell of electrons, then it is more likely to
attract an electron to itself than a large atom
with few electrons.
Inorganic Chemistry
Electronegativity 
Pauling defined electronegativity of an atom as
the tendency of the atom to attract electrons to
itself when combined in a compound.
According to Mulliken, electronegativity is the
average of ionization energy and electron
affinity of an atom.
According to Allred and Rochow,
electronegativity is the attractive force between
a nucleus and an electron at a distance
equivalent to the covalent radius.
Inorganic Chemistry
Electronegativity 
Sanderson electronegativity equalization
This model underlies the concept
of electronegativity equalization, which
suggests that electrons distribute
themselves around a molecule to minimize
or to equalize the Mulliken electronegativity.
Allen electronegativity
Electronegativity is related to the average
energy of the valence electrons in a free
atom.
Inorganic Chemistry
Electronegativity 
Allred and Rochow: Describes  as “a force”
needed to remove an e- from an atom’s ground
state.
AR = 0.744 + 0.359Z*
rcov2
where rcov = covalent radius in Å
Note: The larger the Z*, the more electronegative.
Also, the smaller the rcov, the more electronegative.
Inorganic Chemistry
Electronegativity : Allred and Rochow
F = e2 x Zeff
r2
Where e is the charge on an electron, r is the
covalent radius, and Zeff is the effective nuclear
charge.
The F values are converted to electronegativity
values on the Pauling scale values by using an
empirical relationship:
 = 0.744 + 0.359 x Zeff
r2
Inorganic Chemistry
Inorganic Chemistry
Mulliken Electronegativity
He proposed that the average of the
first ionization energy (EI1) and the electron
affinity (EA) should be a measure of the
tendency of an atom to attract electrons:
EABl = ½ (EA-A + EB-B)
Inorganic Chemistry
Electronegativity Scales
Molecule
Bond
Bond energy
(kJ/mol)
F2
F-F
158
Cl2
Cl-Cl
242
FCl
F-Cl
255
EFCl = ½ (ECl-Cl + EF-F)
= ½ (242 +158)
= 200 kJ/mol
Note: the difference between the observed and
calculated value is 55 kJ/mol.
Inorganic Chemistry
Pauling named this difference the ionic resonance
energy, designated ∆′.
For the reaction A2(g) + B2(g) → 2AB(g),
define ∆′ by:
∆′(A-B) = D(A-B) – [D(A-A) + D(B-B)]
2
∆′(A-B) is related to the electronegativity
difference by:
|A – B| = 0.102√∆′
Inorganic Chemistry
Calculate the ionic resonance energy of HCl,
given that: H = 2.2 and Cl = 3.16.
Solution:
∆′ = {(H – Cl)/0.102}2
= 88.6 kJ/mol
Close agreement with experimental ∆′(HCl) = 92.4 kJ/mol.
Inorganic Chemistry
SAMPLE PROBLEM:
Using the scale χ(F) = 3.98, calculate the
electronegativity of Br. Bond dissociation
energies for Br–Br =193 kJ mol-1, F–F = 155
kJ mol-1 and Br–F in BrF =260 kJ mol-1.
SOLUTION:
Introduction of Inorganic Chemistry
Atomic Polarizability ()
A polarizable atom or ion is one with orbitals that
lie close in energy; large heavy atoms and ions
tend to be highly polarizable.
The polarizability, , of an atom is its ability to be
distorted by an electric field.
An atom or ion (an anion) is highly polarizable if
its electron distribution can be distorted readily, if
unfilled atomic orbitals lie close to the highest–
energy filled orbitals.
Introduction of Inorganic Chemistry
Fajan’s rules
Small, highly charged cations have polarizing
ability.
Large, highly charged anions are easily
polarized.
Cations that do not have a noble-gas e- config
are easily polarized.
NB: Rule 3 is particularly important for d–block
elements
Introduction of Inorganic Chemistry
Example
Which would be the more polarizable, an F- ion or
an I- ion?
F- is small and singly charged
I- ion is identically charged but larger
Therefore: I- ion is likely to be more polarized
EXERCISE!!
i. Account for the decrease in first ionization energy
between fluorine and chlorine.
ii. Suggest a reason why the increase in Z* for a 2p eis smaller between N and O than between C and N.
iii. Account for the large decrease in electron affinity
between Li and Be despite the increase in nuclear
charge.
iv. Which would be more polarizing, Na+ or Cs+?
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