Chemistry

1. Inorganic Chemistry
Structure and Bonding
 Norda Stephenson – norda.stephenson02@uwimona.edu.jm
 Office at the back of C5. Consultation: By appointment
 8 Lectures, 5 Tutorials–
1 – Zeff, Review Periodic Trends, Electronegativity,
2 – Covalent Bonding – Lewis Structures
3- Covalent Bonding – VSEPR
4 – Symmetry
5 – Covalent Bonding – Molecular Orbital Theory
6 - Covalent Bonding – MO Theory
7 – Ionic Bonding
8 – Ionic Bonding
 Final Exam
9 MCQs + 2 structured questions (25 marks each) on Paper 2.
(You must do all MCQs + choose 1 structured.)

2. Tips on how to excel in CHEM1901
 ATTEND ALL YOUR LECTURES AND TUTORIALS
 POSTED LECTURE NOTES SHOULD BE USED ONLY AS A GUIDELINE
WHEN YOU STUDY. THINGS MAY BE SAID AND DONE IN THE
LECTURE ROOM THAT ARE NOT MENTIONED IN THE LECTURE.
 READ ALONG IN YOUR PRESCRIBED TEXT AS WELL AS OTHER
READING MATERIAL. DO SOME ADDITIONAL READING FROM THE
TEXT ON EVERY LECTURE THAT IS GIVEN.
 ATTEMPT ALL TUTORIAL QUESTIONS
 PARTICIPATE IN THE DISCUSSION IN YOUR TUTORIALS
 ENJOY UNIVERSITY LIFE AS YOU LEARN!

3. Why Study Inorganic Chemistry?

4. Why Study Inorganic Chemistry?
Just recall what you’ve done since you opened your eyes this morning…..
 Showered with soaps made with inorganic (metal-containing) catalysts and
used shampoos containing Zn and Se
 Dressed in synthetic fabric made using inorganic catalysts
 Ate breakfast consisting of foods which containing essential trace
metals
 Brushed teeth using toothpaste containing stannous fluoride
 Drove here………….etc.
Point???
Inorganic Chemistry is widely applicable and therefore extremely important
Inorganic chemists are therefore in demand to synthesize or improve new
materials with advantageous properties in all aspects of life. The manipulation of
chemicals in the development of new materials requires knowledge of the
Structure and Bonding in various chemical compounds since this determines
reactivity.

5. Structure and Bonding - Periodic Trends
 Learning Objectives
 1. Discuss trends in atomic radius, ionization energy, electron
affinity, electronegativity and effective nuclear charge.
(REVIEW ON YOUR OWN – Use lecture notes to guide you)
 2. Calculate effective nuclear charge and shielding constant
using Slater’s rules
 3. Calculate electronegativities and bond energies from the
given equation.
 4. Appreciate why it is important to study Inorganic Chemistry
References:
(1) Inorganic Chemistry, Housecroft and Sharpe, Chapter 1 pg. 18-26 and Chapter 2 pg. 42-44.
(2) Basic Inorganic Chemistry, Cotton, Wilkinson and Gaus, 3rd Edition, Chapter 2, pg. 55-66.

6.  With so many metals and non-metals other than carbon (organic chemistry),
with so many variations in oxidation states, coordination numbers, allotropes
etc…….WE NEED A SYSTEM to make sense out of the wide range of
behaviours that exist for the wide range of chemical compounds.
What is this System????
THE PERIODIC TABLE OF ELEMENTS:
A periodic (regular) repetition of chemical and physical properties occur when elements are
arranged in order of increasing atomic number (“The periodic Law”, Mendeleev and
Meyer)

7. Periodic Trends – The Periodic Table of
Elements
This version of the Periodic Table summarizes the orbitals in which valence
electrons reside for each element in the periodic table. Note that valence
electronic structure is the same within a group. Hence?
What in fact does valence electronic structure mean and what is the
consequence of it being the same in a group???????

8. How do we account for the trends that
we have observed with respect to
atomic radius, ionization energy,
electron affinity and electronegativity?
The effective nuclear charge is the main
explanation for the trends that we
observe in atomic radius, ionization
energy, electron affinity and
electronegativity.

9. Effective Nuclear Charge, Z* or Zeff
 The net nuclear charge experienced by electrons in different
atomic orbitals of a multi-electron atom is called the effective
nuclear charge, Z*.
 Z* felt by an electron is affected by the nuclear charge (Z) and
the shielding (screening) effect of other electrons in more
penetrating orbitals.
For a given quantum number, n,
which electrons should experience
the greatest Z*, s or p??????

10. Effective Nuclear Charge, Z* or Zeff
 Z* = Z-σ where Z is the actual number of protons
(atomic number) and σ is the screening or shielding
constant.
What is σ and how do we calculate it?
σ, sigma, (sometimes written as S) is a measure of how well
electrons shield one another from the nuclear charge
 Z* increases continually from left to right on Periodic
Table because of imperfect shielding by electrons in
orbitals of the same principal quantum number. It
however decreases down a group. Why?

11.  Estimating effective nuclear charge and shielding constant based on Slater’s
Rules (one electron model):
Z* = Z-σ
Write out electron configuration in the following groups:
(1s) (2s,2p) (3s,3p) (3d) (4s,4p) (4d) (4f) (5s,5p) etc
Rules for calculating σ for s or p electron
Consider a particular electron in a ns or np orbital:
 1. All electrons in groups higher than that of electron under consideration
each contribute 0 to σ.
 2. Electrons in the same (ns, np) group as the electron under consideration
each contribute 0.35 to σ (except a 1s e- which contributes 0.30 to the
shielding of the other 1s e-).
 3. Electrons in the (n-1) shell each contribute 0.85 to σ.
 4. Electrons in deeper shells each contribute 1.00 to σ.
Effective Nuclear Charge, Z* or Zeff

12. Effective Nuclear Charge
1. In Class Exercises 1 - December 2005 Exam Paper
Question 1(a)i -
Calculate the effective nuclear charge of a 3s electron in Cl. (4 marks)
What would be Z* for a 3p electron in Cl?
Do these estimates of Z* in for 3s and 3p electrons raise
any questions in your mind?
For a given quantum number, n, which electrons
should experience the greatest Z*, s or p??????
Rules for calculating σ for s or p electron
1. All electrons in groups higher than that of electron under
consideration each contribute 0 to σ.
2. Electrons in the same (ns, np) group as the electron under
consideration each contribute 0.35 to σ (except a 1s e- which
contributes 0.30 to the shielding of the other 1s e-).
3. Electrons in the (n-1) shell each contribute 0.85 to σ.
4. Electrons in deeper shells each contribute 1.00 to σ.

13. Effective Nuclear Charge, Z* or Zeff
In Class Exercises 2. Question from Page 21 Inorganic Chemistry 3rd
Edition:
Confirm that the experimentally observed electronic configuration of K,
1s22s22p63s23p64s1 is energetically more stable than the configuration
1s22s22p63s23p63d1.
Rules for calculating σ for d or f electron
(1s) (2s,2p) (3s,3p) (3d) (4s,4p) (4d) (4f) (5s,5p) etc
1. All electrons in levels higher than that of electron under
consideration contribute 0 to σ.
2. Electrons in the same(nd) group each contribute 0.35 to σ.
3. Electrons in all lower groups (including (ns,np)) each
contribute 1.00 to σ.
NOTE: The lower Z* raises orbital energy and makes the electron easier to remove.
Hence when outer electron experiences larger Z*, configuration is more stable.

14. Periodic Trends-Atomic Radius
Atomic radius decreases across a period – Z* increases
Atomic radius increases down a group as n increases: as we move down a
group each member has one more level of inner electrons which effectively
shields outer electrons, the atoms get larger as n increases and Z*
decreases.
van der Waals vs covalent vs ionic radii - to be discussed in tutorial
Other properties such as ionization energies and atomic radii also follow
trends which are affected by electronic structure and Z*.

15. Periodic Trends – Ionisation Energy
The first ionisation energy is the internal energy change at 0K
associated with the removal of the first valence electron from
neutral isolated atoms (gas).

16. Periodic Trends – Ionisation Energy
Trend: General increase across a period .
Rationalize in terms of atomic radius and Z*

17. Periodic Trends – Ionisation Energy
The variation across a period is irregular
Rationalize in terms of Z*.
IONISATION
ENERGY
eV
ATOMIC NUMBER
Li
Be
B
C
N
O
F
Ne
s
1
s
2
s
2
p
1
s
2
p
2
s
2
p
3
s
2
p
4
s
2
p
5
s
2
p
6
E
E
E
+
+ e

18. Class Exercise December 2007 Exam Paper Question
 1a) There is a general increase in ionization energy across
period two of the periodic table but the variation is irregular
with a decrease from Be to B and again from N to O.
i) Account for the general increase in ionization energy across
period two. (3)
i) Using Slater’s rules calculate the effective nuclear charge (Z*)
of a 2s electron in Be and the 2p electron in B. (5)
ii) Suggest why the observed trend in ionization energy (see
above) varies from the trend suggested by Z* for Be and B
calculated from Slater’s rules. (6)
Lesson from this question:
Slater’s rules oversimplify and treat shielding from electrons in s and p
orbitals as the same. However in reality (see radial distribution function) s
orbitals are more penetrating than p and hence shield the p orbitals from the
nuclear charge. ie contribution to  is bigger in reality. However to a large
extent Slater’s rules work as an approximation to Z*.

19. Periodic Trends – Electron Affinity
+Z
Z-
+Z
(Z+1)-
e-
Energy
•Minus (the negative of) the internal energy change associated with
the addition of an electron to a gaseous atom or ion, at 0K
Note the difference between
electron affinity (EA ) and enthalpy
change of electron attachment
(∆EAH) which is sometimes listed.

20. Periodic Trends – Electron Affinity
 Note that the halogens have the highest EA values, and
Cl has the highest value of the halogens
 The addition of a second electron to an element (2nd EA)
is expected to be much less favoured since there will be
repulsion between the negatively charged electron and
the overall negatively charged anion. For example, for O
the values are:
O(g) + e → O-(g) EA = +141 kJmol-1/ ΔHEA= -141kJmol-1
O-(g) + e → O2-(g) EA = - 798 kJmol-1/ ΔHEA= +798kJmol-1

21. Periodic Trends - Electronegativity ()
 The modern definition of  originated with Linus Pauling:
“The power of an atom in a molecule to attract electrons to itself”
 Where ionization energy is high (the electrons are not readily
released) and where electron affinity is high (it is energetically
favourable to acquire electrons), there will be high
electronegativity ().
 first quantified by Pauling in 1932. Now many scales exist
1. Pauling’s Scale
2. Allred and Rochow Scale
3. Mulliken
 Mulliken’s scale of absolute electronegativities is based on
averaging the ionization energy and electron affinity:
 = ½ (IE + EA)

22. Periodic Trends - Electronegativity ()
General increase across a period and decrease
down a group. Rationalize in terms of Z*.

23. Periodic Trends - Electronegativity ()
Pauling’s Scale
Pauling’s scale is based on differences in bond energies and was
developed from bond energy data.
Question: What would you expect the bond energy of H-F to be?
Should it be the average of the bond energies of H-H and F-F?
What is the actual bond energy of H-F, and how do we explain
this?
D(H-F) = 566 kJ/mol D(H-H) = 436 kJ/mol D(F-F) = 158 kJ/mol
How do we explain this?? See below…
 Pauling said ". . . the energy of an actual bond between unlike
(heteronuclear) atoms is greater than the energy of a normal
covalent bond between these atoms.” This additional bond
energy is due to the “additional ionic character of the bond”
called “ionic resonance energy” and it arises because there is
a difference in electronegativities of the two unlike atoms in
the bond which creates an attraction b/w partial charges and
increases the energy required to break the bond.

24. Periodic Trends – Electronegativity
Pauling’s Scale
 Pauling said that the energy of the heteronuclear bond, A-B, can
be considered to be the sum of the nonpolar and polar
contributions i.e.
D(A-B) = Dnp(A-B) + Dp(A-B) (eqn 1)
(where D(A-B) is the bond energy for bond between A and B)
 The nonpolar contributions is the average of the bond energies
for the homonuclear bonds i.e.
 Dnp(A-B) = ½ {Dnp(A-A) + Dnp(B-B)} (eqn 2)
where D(A-A) and D(B-B) are bond energies for A-A and B-B
 The polar contribution, Dp(A-B) is Δ (ionic resonance energy)
Dp(A-B) = Δ = D(A-B) - ½ {D(A-A) + D(B-B)} (eqn 3)

25.  D(A-B) = ½ {D(A-A) + D(B-B)} + Δ (eqn 4)
 And (A - B) = 0.102 (Δ)1/2 (eqn 5)
By assigning the electronegativity of F as 4.00, the
electronegativity of all other elements can be determined
from bond energy data.

26. Class Exercise
Similar to December 1999 Question 1(b)
Using the following data determine H.
Bond Energy HF = 566 kJ/mol
Bond Energy H2 = 436 kJ/mol (remember F = 4.00)
Bond Energy F2 = 158 kJ/mol
(6 marks)

27. Periodic Trends – Electronegativity
Allred and Rochow Scalek
 This scale is based on the coloumbic force of
attraction of the atom for electron density in the
bond
Force = (Z*e)(e)/4εr2
where Z* is the effective nuclear charge, e is the
charge on the electron and r is the mean radius of
the electron, essentially rcov.
 = 0.359Z*/r2 + 0.744
Thus as shown in the Periodic Table, the atoms with
highest electronegativity (eg. fluorine) are those
with smallest atomic radii and largest effective
nuclear charge.

28. Check your understanding
Can you …
 1. discuss trends in atomic radius, ionization energy, electron affinity,
electronegativity and effective nuclear charge?
 2. calculate effective nuclear charge and shielding constant using
Slater’s rules?
 3. calculate electronegativities and bond energies using given equation?
 4. appreciate why it is important to study Inorganic Chemistry?
Assess Yourself : Attempt the Tutorial
Prepare for next class: Read lecture 2 – Practice drawing Lewis structure