ATOMIC STRUCTURE
1
The Atom
1.1
Subatomic Particles
1.2
Atomic Number, Mass Number and Isotopes
1.3
Deducing the Number of Protons, Neutrons and Electrons in an Atom / Ion
1.4
Isotopes
1.5
Determination of Relative Atomic Mass
1.5.1
Mass Spectrometry
1.5.2
Features of a Mass Spectrometer
1.5.3
Relative Abundance (Isotopic Composition and Relative Atomic Mass)
1.5.4
Mass Spectra
1.6
Applications of Mass Spectrometry
2
Electron Configuration
2.1
Relative Energies of Orbitals
2.2
Atomic Orbitals and Electron Density Plots
2.3
Writing Electron Configurations
2.4
Drawing Orbital Diagrams
2.5
Relationship between Electron Configurations and Position in the Periodic Table
3
Ionization
3.1
Ionization Energy
3.2
Patterns in Successive Ionization Energies
3.2.1 Evidence of Energy Levels and Sublevels
3.2.2 Deducing the Group of an Element from Successive Ionization Energy Data
Recall the following definitions:

Relative atomic mass, Ar, of any element is the weighted average mass of its isotopes when
compared to one-twelfth the mass of a carbon-12 atom

Relative molecular mass, Mr, of an element or compound is the mass of one molecule of that
element or compound compared to one-twelfth the mass of a carbon-12 atom
1
The Atom
Page 1 of 30
Dalton:
all matter is made up of individual particles called atoms, which cannot be created or
destroyed; atoms of the same element are alike in every single way
Thomson:
negatively charged electrons scattered in a positively charged sponge-like substance
Rutherford:
the atom is mainly empty space with a small dense positively charged centre, the nucleus
Bohr:
hydrogen pictured as a ‘solar system’, with electrons moving in energy levels around a
positively charged nucleus; presence of neutrons (neutral charge) ensures stability of the
nucleus of elements with more than 1 proton
1.1
Subatomic Particles

Atoms are made up from smaller subatomic particles – protons, neutrons and electrons.

Protons and neutrons are found in the positively charged dense nucleus of the atom.
Collectively, protons and neutrons are called nucleons.

Electrons move about in a volume of space around the nucleus.

This volume of space is also called an electron cloud.

There are discrete energy levels in this volume of space where there is a high
probability of finding an electron.

An electron density map represents the probability of finding an electron in a volume of
space.
Nucleus (protons, neutrons)
Electrons (distributed in region of space around nucleus)
Subatomic
Particle
Symbol
Relative mass
Relative charge
proton
p
1
+1
neutron
n
1
0
electron
e
5  10-4
(considered to be negligible)
–1

The mass of an atom is concentrated in its small, positively charged nucleus.

The electrostatic force of attraction between the protons and electrons holds the atom
together; the neutral neutrons stabilise the nucleus.
Page 2 of 30
1.2
Atomic Number, Mass Number and Isotopes

Atomic number, Z, is the number of protons in the nucleus of an atom. It is the defining
property of an element.

Mass number, A, is the total number of protons and neutrons in the nucleus of an atom.
Mass number
= no. of p + n
A
Z
X
symbol of the element
Atomic number
= no. of p
(= no. of e for atoms)

Isotopes are atoms of the same element which contain the same number of protons but
different number of neutrons (i.e. having the same atomic number but different mass
numbers).
Practice Example 1
Isotopes
12
6
C
13
6
C
14
6
C
n+p
12
13
14
p
6
6
6
n
6
7
8
e
6
6
6
Isotopes
chlorine-35
chlorine-37
Mass number
35
37
Atomic number
17
17
No. of neutrons
18
20
No. of electrons
17
17
Page 3 of 30
1.3

Deducing the Number of Protons, Neutrons and Electrons in an Atom / Ion
The composition of a particular atom or ion can easily be deduced from the proton number,
mass number and charge on the ion.
Consider an atom, AZ X , a positive ion,

A
Z
Xn+ , or a negative ion AZ Xn-
Number of
protons
Number of
neutrons
Number of
electrons
X
Z
A–Z
Z
Symbol
Atom
A
Z
Cation
A
Z
Xn+
Z
A–Z
Z–n
Anion
A
Z
Xn-
Z
A–Z
Z+n

Cations are formed when an atom loses one or more electrons.

Anions are formed when an atom gains one or more electrons.
Practice Example 2
No. of protons (Z)
No. of neutrons (A – Z)
35
17
17
18
37
17
Cl
17
20
16
8
O2-
8
8
31
15
15
16
Species
Cl
P3-
1.4

No. of electrons
17
17
8 + 2 = 10
15 + 3 = 18
Isotopes
Isotopes of the same element have identical chemical properties but different physical
properties.

identical chemical properties: presence of same number of electrons (and hence
valence electrons) results in isotopes undergoing similar chemical bonding and
reactions

different physical properties: different number of neutrons (and hence mass
numbers) results in isotopes having slightly varying density, boiling point, diffusion
rates, with the lighter isotope diffusing more rapidly
Page 4 of 30
Applications of radioisotopes

Radioisotopes contain nuclei that break up spontaneously with the emission of radiation
(which could be alpha particles, beta particles or gamma rays) involving a characteristic
half-life which is the time taken for half of the radioactive nuclei to decay

cobalt-60 in radiotherapy: damage the DNA of cancer cells by using 60Co (which can
undergo beta and gamma decay) so that the cancer cells cannot undergo cell division

iodine-125 as medical tracer: kill brain tumour tissue cells using 125I (which can
undergo beta and gamma decay) so that the tissue cells cannot undergo cell division

iodine-131 as medical tracer: diagnose whether a thyroid gland is functioning normally
using 131I (which can undergo beta and gamma decay) so that the path of the
radioisotope can be traced

carbon-14 in radiocarbon dating:
estimate the age of sample by
comparing the content of 14C (which can
undergo beta decay) in a dead organic
sample with that in living tissue
1.5

Determination of Relative Atomic Mass
The relative atomic mass (Ar) of an element is the weighted average mass of its isotopes
compared to one-twelfth of the mass of carbon-12 atom.
Ar =  relative isotopic mass of each isotope  relative (or percentage) abundance
1.5.1


Mass Spectrometry
A mass spectrometer is used for the accurate determination of the
i)
relative atomic masses of atoms
ii)
relative molecular masses of molecular compounds
iii)
accurate mass of an individual nuclide/ isotope
iv)
identity/ structure of compounds
Advantages of mass spectrometry:

Requires only very little sample (10-12 g)

Accurate

Fast
Page 5 of 30
1.5.2

Features of a Mass Spectrometer
The working of a mass spectrometer may be summarised into 6 key steps:
The sample to
be tested is
vapourized
and injected
into the mass
spectrometer.

The
vapourised
sample is
bombarded
with highenergy
electrons
which
collide with
the atoms of
the sample.
The beam of
positive ions
passes
through a
velocity
selector
which
ensures all
ions have
the same
velocity.
The atoms
lose an
electron to
form mainly
singly
charged
positive
ions (doubly
charged
positive ions
may also be
formed
occasionally).
The angle of deflection, θ depends on the charge-to-mass ratio
)θ



angle of deflection
θ
The ions
are then
accelerated
by an
electric field
into the
magnetic
field, which
causes the
ions to be
deflected
into circular
paths.
q
of the ion.
m
Each circular
path of ions
is brought to
focus onto
the detector
which
detects the
number of
ions passing
through at
each
magnetic
field setting
and
recorded as
peaks in the
form of a
mass
spectrum.
q
m
For ions of the same charge, those with smaller mass are deflected to a greater extent.
e.g. Angle of deflection: 1H+ > 2H+ > 3H+
For ions of the same mass, the more highly charged ions are deflected to a greater extent.
e.g. Angle of deflection: 4He+ < 4He2+
Ions of the same
q
will be deflected to the same extent.
m
e.g. Angle of deflection: (12C1H4)+ =
(14N1H2)+
Page 6 of 30

A mass spectrum is a plot of relative (or percentage) abundance against mass/charge
(m/z ) ratio.
1.5.3
Relative Abundance (Isotopic Composition and Relative Atomic Mass)
Practice Example 3
Chlorine atoms exists as a mixture of 2 isotopes, chlorine-35 and chlorine-37. Given that 25% of
naturally chlorine atoms are chlorine-37 atoms, account for the relative atomic mass of chlorine
that is reflected in the Periodic Table.
% of chlorine-35 = 75 %
Ar of Cl
=
25
100
 37 +
75
100
 71
= 35.5
This value of Ar is comparable to the value reflected in the Periodic Table of 35.45.
Practice Example 4
The relative atomic mass of gallium is 69.7. Gallium is made up of two isotopes
Calculate the percentage abundance of 69Ga.
69
Ga and
71
Ga.
Let the percentage abundance of 69Ga be x%
69.7 =
x
 69 +
 71
= 65
The percentage abundance of 69Ga is 65%
Page 7 of 30
1.5.4

Mass Spectra
The mass spectrum of an element provides the following information:
Feature of Mass Spectrum
Information deduced
number of peaks or lines
number of isotopes present
m/z value of each peak
relative isotopic mass of each isotope
(assume z = 1, i.e. only singly charged ions)
relative (or percentage) abundance of each isotope
Height of each peak

The relative atomic mass, Ar, of an element can be calculated by determining the weighted
average of isotopic masses according to their relative abundance.
Practice Example 5
Deduce the relative atomic mass of boron from the data given in its mass spectrum.
m/z
Species
B+
Isotope
Relative Abundance
10
10
10
23
11
11
11
100
B+
23
B
B
100
Ar of B = ( 123  10) + ( 123  11)
= 10.8
Page 8 of 30
Self Practice 1
Deduce the relative atomic mass of Mg from its mass spectrum.
relative
abundance
8
1
24
m/z
25
26
Species
m/z
Isotope
Relative Abundance
24
24
Mg+
24
8
25
25
Mg+
25
1
26
26
Mg+
26
1
Ar of Mg = (
8
10
 24) + (
1
10
Mg
Mg
Mg
 25) + (
1
10
 26)
= 24.3
Self Practice 2
Deduce the relative atomic mass of the element iron from the data given.
% abundance
91.68%
5.84%
54
m/z
Species
Fe+
2.17%
56
0.31%
57
58
Isotope
m/z
% Abundance
54
54
54
5.84
56
56
56
91.68
57
57
57
2.17
58
58
58
0.31
Fe+
Fe+
Fe+
Ar of Fe =
Fe
Fe
Fe
Fe
(54  5.84) + (56  91.68) + (57  2.17) + (58  0.31)
100
= 55.9
Page 9 of 30

For elements that exist as covalently bonded molecules (eg H2, Cl2, P4, S8):

mass spectrum consists of peaks corresponding to molecules as well as isotopes

molecule will give rise to more than 1 peak if more than 1 isotope is present

relative abundance of these peaks can be calculated from the presence/ combination of
various isotopes
Practice Example 6
The element chlorine has 2 isotopes
spectrum of chlorine gas (Cl2).
35
Cl and
37
Cl of relative abundance 3:1. Deduce the mass
Possible peaks in mass spectrum of Cl2:
m/z
Species
35
(35Cl)+
3
37
(37Cl)+
1
70
(35Cl35Cl)+
9
72
(35Cl37Cl)+
6
74
(37Cl37Cl)+
1
Relative Intensity
Relative abundance
Page 10 of 30
1.6

Applications of Mass Spectrometry
Some applications of mass spectrometry include:

radioactive dating

drug testing (detection of anabolic steroids)

space research

identification of synthesised compounds (in the pharmaceutical industry)
Radioactive Dating

Depends on radioactive carbon-14 which occurs naturally in the atmosphere.

Plants take up CO2 containing 14C (along with 12C) during photosynthesis.

The proportion of 14C : 12C in living matter is exactly the same as in the atmosphere

When an organism dies, it stops taking in carbon. The unstable
radioactive decay over time, causing the 14C : 12C ratio to decrease.

Half life of 14C is 5730 years.

The level of 14C remaining in a sample can be determined by mass spectrometry.

The amount of 14C detected is compared to calibration plots to deduce the age of the sample.
14
C undergoes slow
Drug Testing

Anabolic steroids, which can be used to build bigger muscles and enhance sport
performance, are an artificial form of the male sex hormone testosterone.

Both men and women produce testosterone, together with epitestosterone in their bodies. The
normal ratio of testosterone to epitestosterone (T:E ratio) does not exceed 4:1.

Taking anabolic steroids raises the T:E ratio.

The T:E ratio in urine samples can be determined using mass spectrometry.

When synthetic testosterone breaks down, its products also have a different ratio of
12
C as compared to the breakdown of natural testosterone.

Mass spectrometry can also detect this
present in the sample.
13
C :
13
C:
12
C to determine is synthetic hormone was
Page 11 of 30
Space Research

Mass spectrometers are sent into space to identify the gases present in the atmosphere of
different planets.

Mass spectrometry has also been used to analyse meteorites to provide a better
understanding of the environment and life forms on other planets.

In space shuttles, mass spectrometry is critical in quickly analysing gases in the shuttle to
warn astronauts of any potential problems.

Smaller portable mass spectrometers have also been fitted onto the astronauts’ suits to
detect traces of leaking gases during space walks to serve as safety warnings to astronauts.
2

Electron Configuration
The arrangement of electrons within these energy levels is called its electron configuration and
can be deduced from the atomic number of the atom.
2.1
Relative Energies of Orbitals

Electrons in an atom exist in discrete energy levels in a region of space around the
nucleus. These spaces, each having a characteristic energy level, are known as orbitals.

An orbital can be defined as a region of space around the nucleus where there is 90%
probability of locating the electron.

Some orbitals are close to the nucleus while others are a distance away.

The arrangement of electrons (in their orbitals) in an atom is referred to as its electron
configuration.
Atomic Emission Spectra

When energy is provided to a sample of hydrogen atoms, some of the energy is absorbed and
electrons are excited from a lower energy level to a higher energy level.

The excited state is unstable. The electron will fall back to the lower energy level known as
its ground state.

Emission spectra are produced when photons are emitted from atoms as excited electrons
return to their ground state.

An electron in the excited state returning to ground state emits energy corresponding to a
particular wavelength in relation to the energy level difference.

The emission spectrum is seen as a series of coloured lines at particular wavelengths (the
colour corresponds to the wavelength of radiation emitted).
Page 12 of 30

The spectrum can thus be regarded as a collection of lines due to different electron transitions.

The line emission spectrum of hydrogen provides evidence for the existence of electrons in
discrete energy levels. These fixed energy levels are labelled as n = 1, 2, ..., ∞.
Main Energy Levels

Each main energy level or shell in which electrons are found is assigned a integer number, n,
i.e. n = 1, 2, 3 etc.
n=4
n=3
n=2
n=1

The number n indicates the average distance of the orbitals from the nucleus.

The smaller the value of n,


the closer the electron is to the nucleus

the more strongly the electron is bound to the nucleus.

the lower the energy level of the electron
Each main energy level can hold a different number of electrons.
Maximum number of electrons each main energy level can hold = 2n2
Main energy level
Maximum no. of electrons
n=1
n=2
n=3
n=4
2
8
18
32

Electrons occupy energy levels from the lowest possible energy level (i.e. innermost electron
shell) to the highest possible energy level (i.e. in the outermost electron shell).

An atom in which all electrons are in the lowest possible energy levels is said to be at
ground state.

Valence electrons have the highest energy

The position of an element in Periodic Table in relation to electron arrangement of an atom:

Period number: number of electron shells occupied

Group number: number of valence electrons
Page 13 of 30
Sub-levels (Subshells)

Each main energy level, n, is divided into sub-levels.

The sub-levels are labelled as s, p, d or f.

Generally, the order of energy levels for the sub-levels (within each quantum shell) is:
s<p<d<f.

Sub-levels contain a fixed number of orbitals (regions of space where there is a high
probability of finding an electron).

Each orbital has a defined energy state for a given electronic configuration and chemical
environment and can hold a maximum of 2 electrons of opposite spin.
Subshell
No. of Orbitals
Type of Orbitals
s
1
s
p
3
px, py, pz
d
5
dxy, dxz, dyz, dz2, dx2-y2
f
7
2.2 Atomic Orbitals and Electron Density Plots
s orbitals

The s orbitals are spherical in shape.

s-orbitals of different shells have the same shape but differ in size.

Size of 1s orbital < 2s orbital < 3s orbital

Distance of electrons from the nucleus in 1s orbital < 2s orbital < 3s orbital
Page 14 of 30
p orbitals

The p orbitals have a “dumb-bell” shape.

There are 3 types of p orbitals – px, py and pz with different orientations in space.

The orbitals within a given subshell (e.g. 2px, 2py, 2pz) have the same energy.

p orbitals of different shells have the same shape but differ in size.

Size of 2p orbital < 3p orbital

Distance of electrons from the nucleus in 2p orbital < 3p orbital
Relationship between main energy levels, sub-levels, orbitals and electrons
Main
energy
level, n
No. of
sub-levels
Type of
sub-level
No. of orbitals
within sub-level
No. of electrons
within sub-level
Maximum no. of
electrons in main
energy level (2n2)
1
1
1s
1
2
2
2
2
2s
1
2
8
2p
3
6
3s
1
2
3p
3
6
3d
5
10
4s
1
2
4p
3
6
4d
5
10
4f
7
14
3
4
3
4
18
32
Page 15 of 30
Relative Energies of Orbitals

The relative energies level of the main energy levels and sub-levels in an atom:
Energy level
n=4
4f
4d
4p
Note:
4s subshell has a
lower energy level
than 3d subshell
(when it is not
occupied by
electrons).
3d
n= 3
4s
3p
3s
n=2
2p
2s
n=1
main energy
levels
1s
sub-level
orbitals
Electron Spin

Two electrons in the same orbital must have opposite spins.

This is so that magnetic attraction which results from their opposite spins can counterbalance
the electrical repulsion which results from their identical charges.

The direction of spin of electrons can be indicated by arrows,
and
to represent anti-clockwise spin.
to represent clockwise spin
Note: Electrons can be thought of as spinning on an axis.
Page 16 of 30
2.3 Writing Electron Configurations

The electron configuration of an element refers to the arrangement of electrons of its atoms
in their main energy levels, sub-levels and orbitals.
Notation for writing electron configuration:
2p1
main energy level
number of electrons in sub-level /
subshell
sub-level/
subshell
2.4 Drawing Orbital Diagrams

The arrangement of electrons in a specie can also be represented in an orbital diagrams,
which shows the arrangement of electrons as well as the spin of the electron.
Notation for orbital diagrams to represent electron configuration:
Example: For main energy level n = 3,
3s
3p
sub-level
3d
orbital
main energy level,
n=3
Aufbau Principle

Electrons occupy the lowest energy orbital first before occupying the higher energy orbitals.
Order of filling orbitals
NOTE
 When adding electrons, fill up 4s before 3d.
 When removing electrons, remove from 4s before 3d.
Page 17 of 30
Pauli’s Exclusion Principle

Each orbital can hold a maximum of 2 electrons. The two electrons must be in opposite
spins.
NOT

2 electrons can occupy the same orbital (region of space) despite their mutual repulsion as
they spin in opposite directions.
Hund’s Rule

Orbitals of a sub-level (subshell) must be occupied singly and with parallel spins before
pairing occurs. This is to ensure electrons are as far apart as possible to minimise electronic
repulsion.
NOT
NOT
Note:

When an orbital contains only one electron, the electron is said to be unpaired.

If the above rules are followed, electrons occupy the lowest possible energy and the resultant
arrangement of electrons is known as ground state electronic configuration. (Lowest energy
state)

When one or more electrons absorb energy and are promoted to a higher energy level, the
atom is said to be in an excited state.
Ground State Electronic Configurations of Atoms
Element
Atomic
No.
He
2
Orbital Diagram for Ground State
Electron Configuration
Write the Ground State
Electron Configuration
Full: 1s2
1s
Full: 1s2 2s2 2p2
C
6
1s
2s
2p
Condensed: [He] 2s2 2p2
Ne
Full: 1s2 2s2 2p6
10
1s
2s
2p
Page 18 of 30
Element
Cl
Atomic
No.
17
Orbital Diagram for Ground State
Electron Configuration
1s
2s
2p
3s
Write the Ground State
Electron Configuration
Full: 1s2 2s2 2p6 3s2 3p5
Condensed: [Ne] 3s2 3p5
3p
Ar
18
1s
2s
2p
3s
2s
2p
3s
Full: 1s2 2s2 2p6 3s2 3p6
3p
Ca
20
1s
Full: 1s2 2s2 2p6 3s2 3p6 4s2
Condensed: [Ar] 4s2
3p
1s
V
4s
2s
2p
3s
23
3p
3d
Full: 1s2 2s2 2p6 3s2 3p6 3d3
4s2
Condensed: [Ar] 3d3 4s2
4s
1s

Cr
2s
2p
3s
24
3p
3d
Full: 1s2 2s2 2p6 3s2 3p6 3d5
4s1
Condensed: [Ar] 3d5 4s1
4s
1s

Cu
2s
2p
3s
29
3p
3d
Full: 1s2 2s2 2p6 3s2 3p6 3d10
4s1
Condensed: [Ar] 3d10 4s1
4s
Page 19 of 30
Element
Atomic
No.
Orbital Diagram for Ground State
Electron Configuration
1s
As
2s
2p
3d
4s
Condensed: [Ar] 3d10 4s2
4p3
4p
1s
2s
2p
3s
Full: 1s2 2s2 2p6 3s2 3p6 3d10
4s2 4p4
34
3p
3d
4s
2s
2p
3s
36
3p
4s
Condensed: [Ar] 3d10 4s2
4p4
4p
1s
Kr
Full: 1s2 2s2 2p6 3s2 3p6 3d10
4s2 4p3
33
3p
Se
3s
Write the Ground State
Electron Configuration
3d
Full: 1s2 2s2 2p6 3s2 3p6 3d10
4s2 4p6
4p
Note:

In 3d-block elements, the 3d and 4s orbitals are close in energy. The 4s orbitals are filled first,
followed by the 3d orbitals (with the exception of Cr and Cu).

Cr atom gains extra stability with half-filled 3d subshell.
The configuration 1s2 2s2 2p6 3s2 3p6 3d5 4s1 is more stable than 1s2 2s2 2p6 3s2 3p6 3d44s2.

Cu atom gains extra stability with a fully-filled 3d subshell.
The configuration 1s2 2s2 2p6 3s2 3p6 3d10 4s1 is more stable than 1s2 2s2 2p6 3s2 3p6 3d9 4s2.
Page 20 of 30
Ground State Electronic Configurations of Ions
Orbital Diagram for Ground State
Electronic Configuration
Particle Formula
Write the Ground State
Electronic Configuration
Full: 1s2 2s2 2p5
atom
ion
atom
F
1s
2s
2p
Condensed: [He] 2s2 2p5
Full: 1s2 2s2 2p6
–
F
1s
2s
2p
Condensed: [Ne]
Full: [Ne] 3s2
Mg
1s
2s
2p
3s
Full: 1s2 2s2 2p6
ion
Mg2+
1s
2s
2p
Condensed: [Ne]
1s
atom
2s
2p
3s
Full: 1s2 2s2 2p6 3s2 3p6 3d7
4s2
Co
3p
3d
Condensed: [Ar] 3d7 4s2
4s
ion
Co2+
1s
2s
2p
3s
Full: 1s2 2s2 2p6 3s2 3p6 3d7
4s2
Condensed: [Ar] 3d7
3p
3d
Page 21 of 30
Isoelectronic Species

Atoms/ions with the same electronic configuration (same number of electrons) are said to
be isoelectronic.
Isoelectronic species
Electronic configuration
Note
Ne
O2-
1s2 2s2 2p6
Na+
Ar
P3-
1s2 2s2 2p6 3s2 3p6
Electron configuration can also be
represented using a Noble Gas
‘Core’
e.g. electron configuration of O2represented as [Ne] and that of
Ca2+ as [Ar]
Ca2+
Self Practice 3
For each of the following species,
(a) draw the orbital diagram to show its electron configuration
(b) write its full electron configuration and its condensed form using the noble gas core.
(i)
Ge
orbital diagram:
full electron configuration: 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p2
condensed electron configuration: [Ar] 3d10 4s2 4p2
(ii)
O2–
orbital diagram:
full electron configuration: 1s2 2s2 2p6
condensed electron configuration: [Ne]
(iii)
Fe2+
orbital diagram:
full electron configuration: 1s2 2s2 2p6 3s2 3p6 3d6
condensed electron configuration: [Ar] 3d6
Page 22 of 30
(iv)
Fe3+
orbital diagram:
full electron configuration: 1s2 2s2 2p6 3s2 3p6 3d5
condensed electron configuration: [Ar] 3d5
(v)
Cu+
orbital diagram:
full electron configuration: 1s2 2s2 2p6 3s2 3p6 3d10
condensed electron configuration: [Ar] 3d10
2.5 Relationship between Electron Configurations and Position in the Periodic Table

Elements in the same (main) Groups have the same outermost electronic configuration in their
atoms.
Group
I
II
III
IV
V
VI
VII
0
Outermost
electron
configuration
ns1
ns2
ns2np1
ns2np2
ns2np3
ns2np4
ns2np5
ns2np6
except
He
s-block
d-block
p-block
For main Group elements, No. of valence electrons = Group Number
s-block
p-block
d-block
Page 23 of 30
3
Ionization
Recall: electrons exist in discrete energy levels around the nucleus.

When an electron is at the highest energy level n = , it is no longer in the atom and the atom
is said to be ionized.
3.1 Ionization Energy

The first ionization energy (1st I.E.) of an element is the energy required to remove one mole of
electrons from one mole of gaseous its atoms.
X(g)  X+(g) + eExample
1st I.E. of magnesium:

Mg (g)  Mg+(g) + e
1st I.E. = +736 kJ mol-1
The second ionisation energy (2nd I.E.) of an element is the energy required to remove one
mole of electrons from one mole of singly positively charged gaseous ions.
X+ (g)
 X2+ (g) + e
Example
2nd I.E. of magnesium:
Mg+ (g)  Mg2+ (g) + e
2nd I.E. = +1450 kJ mol-1
Note:

The higher the ionisation energy of an element, the more difficult it is to remove an electron.

The 2nd I.E. > 1st I.E. because more energy is required to remove an electron from a
positive ion (compared to a neutral atom) due to greater net electrostatic attraction.

COMMON MISTAKE:
Mg(g)  Mg2+(g) + 2e represents 2nd ionisation energy !!WRONG!!
The above equation represents the sum of 1St I.E. and 2nd I.E.
3.2 Patterns in Successive Ionization Energies

There is a general increase in successive ionisation energy because an increasing
amount of energy required to remove successive electrons from an increasingly positive
ion due to an increasing NET electrostatic attraction between the nucleus and valence
electrons.
X (g)

X+(g)
+ e
1st I.E.
X+(g)

X2+(g) + e
2nd I.E.
X2+(g) 
X3+(g) + e
3rd I.E.
Page 24 of 30
3.2.1
Evidence of Energy Levels and Sublevels

Successive ionization energy data for an element can show the arrangement of electrons
around the nucleus, i.e. the electron configuration of the element.

The following information can be obtained from a plot of successive ionization energies
against the order of removal of electrons:
(a)
Number of electrons in an atom
(b)
Number of main energy levels occupied and the number of electrons in each
(c)
Number of sub-levels occupied and the number of electrons in each
(d)
Number of valence electrons in the atom
Evidence of Main Energy Levels
The following plot is obtained when all the electrons are successively removed from an atom.
lg I.E.
6
5
4th main energy level
1 electron
2 electrons
1st main energy level
4
3
8 electrons
3rd main energy
level
2
8 electrons
2nd main energy
level
1
0 1
2 3 4 5
6 7 8 9 10 11 12 13 14 15 16 17 18 19
order of electrons removed
Observation from plot
Sharp increase in I.E. for
removal of 2nd electron
Deductions
There is 1 valence
electron
The 2nd electron is in a
different main energy
level
This is the highest-energy
main energy level (furthest
from nucleus)
n=4
Gradual increase in I.E. for
next 8 electrons
There are 8 electrons in
this main energy level
This is the 2nd highest-energy
main energy level (one
inwards from the outermost)
n=3
Another sharp increase in I.E.
for removal of the 10th electron
The 10th electron is
removed from a different
main energy level
This is the 3rd highest-energy
main energy level (two
inwards from the outermost)
n=2
Gradual increase in I.E. for
next 8 electrons
There are 8 electrons in
this main energy level
n=2
Page 25 of 30
Observation from plot
Deductions
The 18th electron is
removed from a different
main energy level
Another sharp increase in I.E.
for removal of the 18th electron
This is the lowest-energy
main energy level (closest to
nucleus)
n=1
There are 2 electrons in
this main energy level
Electron configuration of the atom may be deduced to be: 1s22s22p63s23p64s1
 This is a Group 1 element
Note: Some graphs may not show the successive removal of all electrons.
 In these cases, electron configuration of the element cannot be determined but Group
number of the element can still be deduced.
Evidence of Sub-levels in Main Energy Levels
lg I.E.
3s
3p
1
2
3
4
5
6
7
8
9
order of electrons removed
Zooming in on the section corresponding to the 3rd main energy level in the previous graph, it is
observed that there is

a steady increase in I.E.s for the removal of 2nd to 7th electrons
 these 6 electrons are in the same sub-level

a minor increase (jump) in I.E. for the removal of the 8th electron
 8th electron is removed from a different sub-level which is closer to the nucleus.
There is stronger electrostatic force of attraction between the 8th electron and nucleus
(since it is closer to the nucleus) and more energy is required to remove it.

The above shows that the main energy level is further divided into sub-levels (labelled as
3s and 3p) which are occupied by two and six electrons respectively.
Page 26 of 30
3.2.2
Deducing the Group of an Element from Successive Ionization Energy Data

To determine the Group to which an element belongs to, we must determine the number of
valence electrons.

The point where a big increase in I.E. occurs indicates the removal of successive electrons
from an inner main energy level.
Practice Example 7
From the following data, predict the Group to which this element belongs and write its electron
configuration.
I.E. kJ mol-1
1st
2nd
3rd
4th
900
1758
14905
21060
Work out the difference between the 1st & 2nd IE, 2nd & 3rd IE, 3rd & 4th IE.


There is a sharp increase in I.E. for removal of the 3rd electron
(as indicated by the largest difference in IE between 2nd IE & 3rd IE).

3rd electron is removed from an inner main energy level which is closer to the
nucleus. Hence, there is stronger electrostatic force of attraction between the 3rd
electron and nucleus and more energy is required to remove it

The element has 2 valence electrons

The element belongs to Group 2.
Hence, the valence electron configuration of the element is ns2.
Page 27 of 30
Practice Example 8
The following element has the following successive ionisation energies when all its electrons are
successively removed.
(i)
Deduce the Group number of the
element.
(ii)
Deduce the total number of quantum
shells.
(iii)
Write the electron configuration of the
element.
(i)
There is a sharp increase in I.E. for removal of the 3rd electron
(as indicated by the largest difference in IE between 2nd IE & 3rd IE).
(ii)

3rd electron is removed from an inner main energy level which is closer to the
nucleus. Hence, there is a stronger electrostatic force of attraction between the
3rd electron and nucleus and more energy is required to remove it.

The element has 2 valence electrons

The element belongs to Group 2.
From the graph,
Total number of electrons = 12

There should be 3 main energy levels / 3 electron shells ( i.e. n = 3 )
(Note also there are 3 regions of gradual increase in I.E.)
(iii)
Electron configuration: 1s22s22p63s2
Challenge yourself to explain this: Group I elements have high 2nd I.E.
Group II elements have high 3rd I.E.
Group III elements have high 4th I.E.
Group IV elements have high 5th I.E.
Group V elements have high 6th I.E.
Group VI elements have high 7th I.E.
Group VII elements have high 8th I.E.
Page 28 of 30
Self Practice 4
Deuce the electron configuration of the element from its successive ionization energies given.
Successive removal of 20 electrons from element
Removal of 1st electron: Ca(g) → Ca+(g) + eRemoval of 2nd electron: Ca+(g) → Ca2+(g) + e......
Removal of 19th electron: Ca18+(g) → Ca19+(g) + eRemoval of 20th electron: Ca19+(g) → Ca20+(g) + e-
No. of electrons removed

There is a sharp increase in I.E. for removal of the 3rd , 11th and 19th electrons

3rd , 11th and 19th electrons are removed from main energy levels successively closer
to the nucleus.
There are stronger electrostatic forces of attraction between the nucleus and the
electrons closer to it and more energy is required to remove these electrons

1st and 2nd electrons removed are from the outermost electron shell (highest energy
level) (The element has 2 valence electrons)

3rd to 10th electrons removed are from an inner main energy level (one inwards from
the outermost)

11th to 18th electrons removed are from an inner main energy level closer to nucleus

19th and 20th electrons removed are from the innermost energy level

electron configuration: 1s2 2s2 2p6 3s2 3p6 4s2
Page 29 of 30
Self Practice 5
Deduce the group number for an unknown element given its first eight successive ionization
energies in kJ mol-1.
580

1800
2750
11580
14850
18400
23300
27500
There is a sharp increase in I.E. for removal of the 4th electron

4th electron is removed from an inner main energy level which is closer to the
nucleus.
Hence, there is stronger electrostatic force of attraction between the 4th electron and
nucleus and more energy is required to remove it

1st to 3rd electrons removed are from the outermost energy level

4th to 8th electrons removed are from an inner shell (one inwards from the outermost)

The element has 3 valence electrons

The element belongs to Group 13.
(valence electron configuration of the element is ns2 np1)
Page 30 of 30