CHEM 1103 Atomic Structure 2 form the UOM

Electrons are arranged in shells which are subdivided into subshells and orbitals. They are arranged in terms of energy.
Videos Shells, subshells, and orbitals (video) | Khan Academy
Introduction to electron configurations (video) | Khan Academy
Valence electrons (video) | Khan Academy

Example Aluminium
Electron arrangement in aluminium = 2, 8, 3
The first two electrons are in shell n = 1 which contains 1s sub shell
The second shell n = 2 contains 2s and three 2p orbitals of the same energy
• Aufbau principle - electrons fill the lowest energy orbitals first, and then move
up to higher energy orbitals only after the lower energy orbitals are full.
• Hund's rule states that:
Every orbital in a sublevel is singly occupied before any orbital is doubly
occupied.
All of the electrons in singly occupied orbitals have the same spin (to
maximize total spin).
• Pauli exclusion principle
An orbital cannot have both the electrons in the same spin motion
electrons will be in either positive half spin (+1/2) or negative half spin (-1/2)
1s
2s
2p
The last three electrons are in shell n = 3 which contains 3s, three 3p and five 3d orbitals
3s
3p

Chlorine 17 electrons
1s2 2s22p6 3s23p5
2s
2p
3s
3p
1s
n = 2
n = 1
n = 3

2s2
2p6
3s2
3p6
1s2
n = 2
n = 1
3d2
Titanium 22 electrons
4s2
E
N
E
R
G
Y
Ti - 1s2 2s2 2p6 3s2 3p6 4s2 3d2

Introduction to new quantum numbers
• An atomic orbital is defined by a unique set of three quantum
numbers: n, l and ml.
• An electron in an atom is defined by a unique set of four quantum
numbers: n, l, ml and ms.
• n = principal quantum number which may have values 1, 2, 3, …
corresponding to the first, second, third, …shell of electrons around
the nucleus.
• l = azimuthal (subsidiary) quantum number which may have values 0,
1, 2, …(n-1). Describes the orbital angular momentum or shape of the
orbital.
n = 4, l = 0, 1, 2, 3
l = 0 (s orbital); l = 1 (p orbital); l = 2 (d orbital); l = 3 (f orbital)

Introduction to new quantum numbers
• n = principal quantum number which may have values 1, 2, 3, …
corresponding to the first, second, third, …shell of electrons around
the nucleus.
• l = azimuthal (subsidiary) quantum number which may have values 0,
1, 2, …(n-1). Describes the orbital angular momentum or shape of the
orbital.
n = 4, l = 0, 1, 2, 3
l = 0 (s orbital); l = 1 (p orbital); l = 2 (d orbital); l = 3 (f orbital)

Introduction
• m or ml = magnetic quantum number which may have values
from +l, l-1, …..0……-l (2l+1) values. Represents the
component of the orbital angular momentum along the z-axis
(orientation).
If l = 2, m can take 5 values (2, 1, 0, -1 and -2) representing the
5 d orbitals.
• ms = magnetic spin quantum number which has a value of either
+½ or -½. Represents the component of the spin quantum number
lying parallel or antiparallel to the magnetic field.

Electronic configuration of atoms
• The arrangement of e in atoms is known as the electronic configuration.
• Electrons are arranged in orbitals according to a set of rules (Aufbau
Principle):
 Each orbital can hold a maximum of two electrons.
 Electrons occupy the orbital of lowest energy first.
 Hund’s rule: when several orbitals have the same energy, they are not
paired if this can be avoided. Thus, in the ground state, an atom will
contain the maximum number of unpaired electron spins.
 The Pauli exclusion principle: No two e in one atom can have all four
quantum numbers the same. Two e in the same orbital have opposite
spins.

Each orbital can be represented by a square box.
Ex: Boron (Z=5) 1s2 2s2 2p1
Ex: Carbon (Z = 6) 1s2 2s2 2p2
Valence e
Valence e

(l = 0)
(l = 1)
Then ml = -1, 0, 1
Second shell n = 2 +1/2 -1/2
No electrons have the same set of quantum numbers

All three subshells of the
third shell (3s, 3p and 3d)
have higher energy than
the subshells of the second
shell. However the 4s
subshell of the fourth shell
has lower energy than the
3d subshell.
Energy of orbitals

The atomic configuration of the first row transition series is Sc to Zn (d1 to d10):
1s22s22p63s23p64s23dn

Anomalous Electronic Configurations
• A few exceptions to the Aufbau principles exist.
Examples:
• half-filled d shell:
• Cr has [Ar]4s13d5; Cr has 24 electrons
• Mo has [Kr] 5s14d5
• filled d subshell:
• Cu has [Ar]4s13d10
• Ag has [Kr]5s14d10.
• Au has [Xe]6s14f145d10
• Exceptions occur with larger elements where orbital energies are similar.

All s orbitals are spheres. They only differ in size. The 3s
orbital is bigger than the 2s, which is bigger than the 1s.
Shape of s orbitals

Shapes of p orbitals
The shape of p orbitals is different from the shape of s orbitals. The
three p orbitals of a shell have the same shape, but have different
orientations in space (oriented along the x, y and z axis).
The 3p orbitals have the same shape as the 2p orbitals, but are larger.

Periodic Table of Elements
Divided into Groups
and Periods

Trends in I.E.
O(g)  O+(g) + e ΔHIE1 = +1310 kJ mol-1 (First I.E.)
The electron is removed from the outer subshell of the oxygen atom (i.e. a
2p electron).
The energy required to remove the second electron is called the second
ionization energy.
Second I.E of oxygen:
O+(g)  O2+(g) + e ΔHIE2 = +3400 kJ mol-1
The second I.E. is larger than the first I.E. because more energy is required
to pull an electron away from a positive ion (compared with a neutral
atom).

Factors influencing the I.E. of elements
The I.E. of an atom is strongly influenced by three atomic parameters.
(a) The distance of the outermost electron from the nucleus (size of atom)
The bigger the atom is, the further the electron is from the nucleus.
As this distance increases, the attraction of the positive nucleus for the
negative electron decreases and consequently the I.E. decreases.
(b) The size of the positive nuclear charge
The bigger the nuclear charge, the bigger the attractive force of the
nucleus on the electrons  the larger the energy required to remove the
electron. As the nuclear charge becomes more positive with increasing
atomic number, its attraction for the outermost electron increases and
consequently the I.E. increases.

Factors influencing the I.E. of elements
(c) The screening (shielding) effect of inner shells of electrons
Electrons in inner shells exert a repelling effect on electrons in the
outermost shell of an atom. The outermost electron is screened or
shielded from the attraction of the positive nucleus by the repelling
effect of inner electrons. This screening effect means that the effective
nuclear charge is much less than the full positive charge in the nucleus.
In general the screening effect by inner electrons is more effective the
closer these inner electrons are to the nucleus. Thus, electrons in shells
of lower principal quantum number are more effective shields than
those in shells of higher quantum number.

Calculating the screening effect S
Slater’s rules (Z* = Z-S)
• Effects of screening are generally expressed in terms of an effective
nuclear charge, Z*, which can be viewed as the average value of the
nuclear charge that an electron in an orbital experiences after the
screening effects of other electrons have been taken into account.
• A set of empirical rules - devised by Slater (1932) to calculate Z* or Zeff.
 First write the electronic configuration of the atom/ion in the
following form:
[1s] [2s2p] [3s3p] [3d] [4s4p] [4d] [4f] [5s5p] [5d] [5f]
ns and np electrons always considered as a single group.

Slater’s rules (Finding S)
(1) For an e in an [ns, np] group, e in groups to the right contribute nothing
to shielding.
(2) For an e in an [ns, np] group, other e in the same group contribute 0.35
charge units to the shielding, except for 1s electrons, which contribute
0.3. example [2s2 2p1] S = 2 x 0.35
(3) Each e in the (n-1) group contributes 0.85 to the shielding.
(4)Each e in the (n-2) or lower group contributes 1.0 to the shielding.
(5) For an e in an nd or nf group, rules (1) and (2) remain the same, and all
e in groups to the left contribute 1.0 to the shielding.

Slater’s Rules for Calculating Zeff
• Write out the electron configuration in groups using the following
order:
(1s) (2s, 2p) (3s, 3p) (3d) (4s, 4p) (4d) (4f) (5s, 5p) …
• Identify the group in which the electron of interest lies; ignore
electrons to the right of this group
• If the electron of interest is an s or p electron, then each additional
electron in its (ns, np) group contributes 0.35 to s, each electron in
the n – 1 shell contributes 0.85 to s, and each electron further to the
left contributes 1.00 to s.
• If the electron of interest is a d or f electron, then each additional
electron in its (nd) or (nf) group contributes 0.35 to s and each
electron further to the left contributes 1.00 to s.

Slater’s rules
Example - Determine Zeff = Z – S felt by the outer 4s electron for
potassium atom, Z=19.
[1s2] [2s22p6] [3s23p6] 4s1
S = 10 + (0.85 x 8) = 16.8 Zeff = 19 – 16.8 = 2.2
each electron in the n – 1 shell contributes 0.85 to s, and each electron
further to the left contributes 1.00 to s.

Slater’s rules
• Calculate the Zeff felt by (a) one of the 4s electrons and (b) one of
the 3d electrons in vanadium (Z=23)
(Ans: 3.3 and 4.3)

Examples
Calculate the effective nuclear charge felt by a
(i) 4s electron;
(ii) 3d electron;
(iii) 3s or 3p electron;
(iv) 2s or 2p electron;
(v)1s electron in the iron atom (Z = 26).
Ans: (i) 3.75; (ii) 6.25; (iii) 14.75; (iv) 21.85; (v) 25.70

Trend in I.E. across a period
Moving from left to right across any period there is a
general increase in the first I.E. because:
 The nuclear charge is increasing. The electrons in the
outermost shell are more strongly bound to the nucleus
due to the increasing effective nuclear charge.
 The distance of the outermost electron from the nucleus is
decreasing because the increasing nuclear charge is
pulling electrons closer to the nucleus, so the atomic
radius is decreasing.

• Although there is a general increase in the first I.E. across any period,
there are breaks in the overall pattern, e.g. the first I.E. of Be is higher
than that of B.
• The electron configuration of Be is 1s2, 2s2, whereas that of boron is 1s2,
2s2, 2p1. All the sub-shells in Be are filled, but the outer 2p sub-shell of
B contains only one electron. There is extra stability associated with
filled sub-shells. The electron removed from p-orbital is farther from the
nucleus and there is a small amount of repulsion by the s electrons.
• In period 2, there is another break with nitrogen which has a higher I.E.
than O. The electronic configurations of N and O are 1s2, 2s2, 2p3 and
1s2, 2s2, 2p4 respectively. Stability associated with a half-filled sub-shell.
Repulsion from other electron in the doubly occupied orbital helps in its
removal.

Trends in First Ionization Energies

Trend in I.E. down a group
Moving down any group in the periodic table, there is a general decrease in
the first I.E. as the proton numbers increase.
Group 2 Atomic No First I.E.
/kJ mol-1
Group 17 Atomic No First I.E.
/kJ mol-1
Be 4 900 F 9 1680
Mg 12 736 Cl 17 1260
Ca 20 590 Br 35 1140
Sr 38 548 I 53 1010
Ba 56 502
Distance and screening factors outweigh the nuclear charge factor and
cause the first I.E. of elts to decrease down a group in the periodic table.

The trend across from left to right is accounted for by a) the
increasing nuclear charge.

Periodicity
•When elements are arranged in order of increasing
atomic number, certain sets of properties recur
periodically.
Metallic, non-metallic and metalloid properties
Atomic radius
Ionization energies
Electron affinities
Electronegativity

Periodicity
•Most obvious trend is from metallic to non-metallic.
•Acid-base properties of oxides:
•Metal oxides tend to be basic
•(Na2O = NaOH in H2O)
•Non-metal oxides tend to be acidic
•(SO2 = H2SO3 in H2O)
•Elements in the same group generally have similar chemical
properties because they have the same number of electrons
in their outer energy levels.

Atomic Radius
• Atomic radii decrease
across a row in the periodic
table due to an increase in
the effective nuclear charge
(increase in the number of
protons)
• Within each group (vertical
column), the atomic radius
tends to increase with the
period number.
Atomic Radii for Main Group Elements
Zeff = Z- S

Ionic radii
•If a neutral atom (e.g. Na) loses an electron, it becomes
positively charged (Na+) and its radius decreases. Cations
are smaller than their parent atoms because the outermost
electron is removed and repulsions are reduced.
•If a neutral atom (e.g. Cl) gains an electron, it becomes
negatively charged (Cl-) and its radius increases. Anions
are larger than their parent atoms as electrons are added
and repulsions are increased.

Sizes of Ions
Ionic size depends
upon:
Nuclear charge.
Number of electrons.
Orbitals in which
electrons reside.
• Ions increase in size
as we go down a
column due to
increasing value of n.

Sizes of Ions
• In an isoelectronic series, ions have the same number of
electrons.
• Ionic size decreases with an increasing nuclear charge.

• Arrange the following in order of size, smallest to largest:
Na, Na+, Mg, Mg2+, Al, Al3+, S, S2-, Cl, Cl-
Ans: Al3+, Mg2+, Na+, Cl, S, Al, Mg, Na, Cl-, S2-
• Predict which of the following substances has the largest
radius: P3, S2, Cl, Ar, K+, Ca2+.

• Energy change accompanying addition of electron to a gaseous
atom
E.g. Cl(g) + e  Cl(g) Eea = 348.6 kJ/mol
• Can be either endothermic or exothermic depending on the
element
• The greater the negative value of the electron affinity, the
greater the tendency of an atom to accept an electron
• A positive value indicates that energy must be absorbed for an
atom to gain an electron.
Electron Affinity

Trends in Electron Affinity
In general, electron affinity becomes more exothermic as you go
from left to right across a row.

• Going from left to right on the periodic chart, general increasing
tendency to form negative ions. However, there are more
exceptions than with ionization energy.
• Electron affinities generally become smaller as we go down a
column of the periodic table for two reasons.
 First, the electron being added to the atom is placed in larger
orbitals, where it spends less time near the nucleus of the atom.
 Second, the number of electrons on an atom increases as we go
down a column, so the force of repulsion between the electron
being added and the electrons already present on a neutral atom
becomes larger.

• From the data in the table the halogens clearly have a strong tendency to
become negatively charged.
• Inert gases and group I & II elements have a very small Eea.
• There are two discontinuities.
• The first occurs between Groups IA
and IIA. Added electron must go in p-
orbital, not s-orbital. Electron is
farther from nucleus and there are
repulsions from s-electrons.
• The second occurs between Groups
IVA and VA. Group VA has no empty
orbitals. Extra electron must go into
occupied orbital, creating repulsion.

ELECTRONEGATIVITY
•Electronegativity is a measure of the tendency of an
atom to attract a bonding pair of electrons.
•The Pauling scale is the most commonly used.
Fluorine (the most electronegative element) is
assigned a value of 4.0, and values range down to
caesium and francium which are the least
electronegative at 0.7.

CHEM 1103 Atomic Structure 2 form the UOM

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