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AQA A Level Chemistry Your notes

Atomic Structure
Contents
Fundamental Particles
Mass Number & Isotopes
Time of Flight Mass Spectrometry
Shells & Orbitals
Electron Configuration
Ionisation Energy
Ionisation Energy: Trends & Evidence

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Fundamental Particles
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Structure of an Atom
All matter is composed of atoms
Atoms are the smallest parts of an element that can take place in chemical reactions
Atoms are mostly made up of empty space around a very small, dense nucleus that
contains protons and neutrons
Protons and neutrons are sometimes referred to as nucleons because they are found in the nucleus
The nucleus has an overall positive charge
This is because the protons have a positive charge and the neutrons have a neutral charge
Negatively charged electrons are found in orbitals in the empty space around the nucleus

The basic structure of an atom (not to scale)

Subatomic Particles

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Subatomic particles are the particles an element is made up of and include protons, neutrons and
electrons
Your notes
These subatomic particles are so small that it is not possible to measure their masses and charges
using conventional units (such as grams and coulombs)
Instead, their masses and charges are compared to each other using ‘relative atomic masses’ and
‘relative atomic charges’
These are not actual charges and masses but they are charges and masses of particles relative to each
other
Protons and neutrons have a very similar mass so each is assigned a relative mass of 1 whereas
electrons are 1836 times smaller than a proton and neutron
Protons are positively charged, electrons negatively charged and neutrons are neutral
The relative mass and charge of the subatomic particles are:
Relative mass & charge of subatomic particles table

Sub-atomic particle Relative electrical charge Relative Mass

Proton +1 1

Neutron 0 (neutral) 1

Electron -1 1
1836

Examiner Tips and Tricks

The relative mass of an electron is almost negligible.
The charge of a single electron is -1.602 x 10-19 coulombs whereas the charge of a proton is +1.602 x
10-19 coulombs, however, relative to each other, their charges are -1 and +1 respectively.

Atoms: Key Terms

The atomic number (or proton number) is the number of protons in the nucleus of an atom and has
symbol Z

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The atomic number is equal to the number of electrons present in a neutral atom of an element
Eg. the atomic number of lithium is 3 which indicates that the neutral lithium atom has 3 protons Your notes
and 3 electrons
The mass number (or nucleon number) is the total number of protons and neutrons in the nucleus of an
atom and has symbol A
The number of neutrons can be calculated by:
Number of neutrons = mass number - atomic number
Protons and neutrons are also called nucleons

Examiner Tips and Tricks

The mass (nucleon) and atomic (proton) number are given for each element in the Periodic Table

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Mass Number & Isotopes

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Subatomic Structure of Atoms & Ions
The mass of an atom is concentrated in the nucleus, because the nucleus contains the heaviest
subatomic particles (the neutrons and protons)
The mass of the electron is negligible
The nucleus is also positively charged due to the protons
Electrons orbit the nucleus of the atom, contributing very little to its overall mass, but creating a ‘cloud’
of negative charge
The electrostatic attraction between the positive nucleus and negatively charged electrons orbiting
around it is what holds an atom together

The mass of the atom is concentrated in the positively charged nucleus which is attracted to the
negatively charged electrons orbiting around it
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An atom is neutral and has no overall charge

Ions on the other hand are formed when atoms either gain or lose electrons, causing them to become Your notes
charged
The number of subatomic particles in atoms and ions can be determined given their atomic (proton)
number, mass (nucleon) number and charge

Protons
The atomic number of an atom and ion determines which element it is
Therefore, all atoms and ions of the same element have the same number of protons (atomic number)
in the nucleus
E.g. lithium has an atomic number of 3 (three protons) whereas beryllium has atomic number of 4 (4
protons)
The number of protons equals the atomic (proton) number
The number of protons of an unknown element can be calculated by using its mass number and
number of neutrons:
Mass number = number of protons + number of neutrons
Number of protons = mass number - number of neutrons

Worked Example
Determine the number of protons of the following ions and atoms:
1. Mg2+ ion
2. Carbon atom
3. An unknown atom of element X with mass number 63 and 34 neutrons
Answers:
1. The atomic number of a magnesium atom is 12 indicating that the number of protons in the
magnesium element is 12
Therefore the number of protons in a Mg2+ ion is also 12
2. The atomic number of a carbon atom is 6 indicating that a carbon atom has 6 protons in its
nucleus
3. Use the formula to calculate the number of protons
Number of protons = mass number - number of neutrons
Number of protons = 63 - 34

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Number of protons = 29
Element X is therefore copper
Your notes
Electrons
An atom is neutral and therefore has the same number of protons and electrons
Ions have a different number of electrons to their atomic number depending on their charge
A positively charged ion has lost electrons and therefore has fewer electrons than protons
A negatively charged ion has gained electrons and therefore has more electrons than protons

Worked Example
Determine the number of electrons of the following ions and atoms:
1. Mg2+ ion
2. Carbon atom
3. An unknown atom of element X with mass number 63 and 34 neutrons
Answers:
1. The atomic number of a magnesium atom is 12 suggesting that the number of protons in the
neutral magnesium atom is 12
However, the 2+ charge in Mg2+ ion suggests it has lost two electrons
It only has 10 electrons left now
2. The atomic number of a carbon atom is 6 suggesting that the neutral carbon atom has 6
electrons orbiting around the nucleus
3. The number of protons of element X can be calculated by:
Number of protons = mass number - number of neutrons
Number of protons = 63 - 34
Number of protons = 29
The neutral atom of element X therefore also has 29 electrons

Neutrons
The mass and atomic numbers can be used to find the number of neutrons in ions and atoms:
Number of neutrons = mass number (A) - number of protons (Z)

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Worked Example
Determine the number of neutrons of the following ions and atoms: Your notes
1. Mg2+ ion
2. Carbon atom
3. An unknown atom of element X with mass number 63 and 29 protons
Answers:
1. The atomic number of a magnesium atom is 12 and its mass number is 24
Number of neutrons = mass number (A) - number of protons (Z)
Number of neutrons = 24 - 12
Number of neutrons = 12
The Mg2+ ion has 12 neutrons in its nucleus
2. The atomic number of a carbon atom is 6 and its mass number is 12
Number of neutrons = mass number (A) - number of protons (Z)
Number of neutrons = 12 - 6
Number of neutrons = 6
The carbon atom has 6 neutrons in its nucleus
3. The atomic number of an element X atom is 29 and its mass number is 63
Number of neutrons = mass number (A) - number of protons (Z)
Number of neutrons = 63 - 29
Number of neutrons = 34
The neutral atom of element X has 34 neutrons in its nucleus

Isotopes
What are isotopes?
Isotopes are atoms of the same element that contain the same number of protons and electrons but a
different number of neutrons
The symbol for an isotope is the chemical symbol (or word) followed by a dash and then
the mass number
E.g. carbon-12 and carbon-14 are isotopes of carbon containing 6 and 8 neutrons respectively

Isotopes of hydrogen

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Your notes

Isotopes of the same element have different numbers of neutrons

RAM from Mass Spectra

Isotopes have the same chemical properties but different physical properties

Chemical properties
Isotopes of the same element display the same chemical characteristics
This is because they have the same number of electrons in their outer shells
Electrons take part in chemical reactions and therefore determine the chemistry of an atom

Physical properties
The only difference between isotopes is the number of neutrons
Since these are neutral subatomic particles, they only add mass to the atom
As a result of this, isotopes have different physical properties such as small differences in their mass
and density
Isotopes are different atoms of the same element that contain the same number of protons and
electrons but a different number of neutrons.

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These are atoms of the same elements but with different mass numbers
Because of this, the mass of an element is given as relative atomic mass (Ar) by using the average
mass of the isotopes Your notes
The relative atomic mass of an element can be calculated by using the relative abundance values
The relative abundance of an isotope is either given or can be read off the mass spectrum
Ar =
( relative abundance isotope 1 × mass isotope 1 ) + relative abundance
(
isotope 2 × mass isotope 2 ) etc
100

Worked Example
Calculating relative atomic mass of oxygen
A sample of oxygen contains the following isotopes:

Isotope Percentage abundance

16O 99.76
17O 0.04
18O 0.20

What is the relative atomic mass of oxygen in this sample, to 2dp?

1. 16.00
2. 17.18
3. 16.09
4. 17.00
Answer:

The correct answer option is A

( 99 . 76 × 16 ) + ( 0. 04 × 17 ) + ( 0. 20 × 18 )
Ar =
100
Ar = 16.0044
Ar = 16.00

Worked Example
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Calculating relative atomic mass of boron

Calculate the relative atomic mass of boron using its mass spectrum, to 1dp: Your notes
Answer:

(19 . 9 × 10) + (80 . 1 × 11 )

Ar = = 10.801 = 10.8
100

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Time of Flight Mass Spectrometry

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Time of Flight Mass Spectrometry
Mass Spectrometry is a powerful analytical technique
It is the most useful instrument for accurate determination of the relative atomic mass of an
element, based on the abundance and mass of each of its isotopes
It is also used to find the relative molecular mass of molecules
As a sample passes through the mass spectrometer, a spectrum is produced of mass / charge ratio
against abundance
The spectrum can be used to find the relative isotopic abundance, atomic and molecular mass and the
structure of a compound
The peak with the highest mass is the molecular ion peak, M+, and the peak which has the largest
abundance (tallest peak) is called the base peak
There are several types of mass spectrometer, but all of them are based on an ionised sample being
accelerated through the mass spectrum, and being separated based on the ratio of their charge to
their mass

Time of Flight Mass Spectrometry

This is a common form of mass spectrometry, where all particles of the sample to be analysed are
ionised to form 1+ ions
These 1+ ions are then accelerated to high speeds, deflected through the spectrometer and then arrive
at the detector heavy and light ions are deflected
As they hit the detector, the mass spectrum graph is produced
The whole of the apparatus is kept under a high vacuum to prevent any ions that are produced from
colliding with molecules in the air

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Your notes

to form + ions

Inside the time of flight mass spectrometer

There are 4 key stages in time of flight mass spectrometry:
Ionisation
Acceleration
Ion drift
Detection

Stage 1: Ionisation
There are two key ways in which the sample could be ionised:
Electron Impact (or electron ionisation)
Electrospray Ionisation
Electron Impact Ionisation
This method of ionisation is used for elements and substances which have a lower molecular mass
The sample is vaporised and then bombarded with high energy electrons
dable leho
The electrons are 'fired' from an electron gun
The electron gun is a hot wire filament which emits electrons as a current runs through it

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As the sample is bombarded by these electrons, an electron is knocked off each particle, forming a 1+
ion
X (g) → X+ (g) + e- Your notes
is x a metal or non - metal
The 1+ ions which have been formed are called molecular ions, or M+ ions
These are then attracted towards a negatively charged plate
This accelerates them through the mass spectrometer
The molecular ion can be broken down further, or fragmented
The fragments are also accelerated through the sample and hit the detector, causing different
peaks to show on the mass spectrum which is produced
Electrospray Ionisation
This method is used for substances which have a higher molecular mass
Unlike with electron impact ionisation, fragmentation is unlikely to happen
This is often called a soft ionisation technique
For this method, the sample is dissolved in a volatile solvent
The solvent is injected into the mass spectrometer using a hypodermic needle
This produces a fine mist or aerosol
The needle is attached to a high voltage power supply, so as the sample is injected, the particles are
ionised by gaining a proton from the solvent
X (g) + H+ → XH+ (g)
The solvent evaporates and the XH+ ions are attracted towards a negatively charged plate
This accelerates them through the mass spectrometer

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Your notes

Time of Flight Mass Spectrometer

Stage 2: Acceleration
The 1+ ions formed from either ionisation method are accelerated using an electric field
They are all accelerated to have the same kinetic energy
This is important for you to remember when completing calculations
Since all 1+ ions will have the same kinetic energy, their velocity will depend on their mass
Lighter ions will move faster and heavier ions will move slower

Stage 3: Ion Drift (in the flight tube)

The 1+ ions will pass through a hole in the negatively charged plate and move into a flight tube
This is where the name 'Time of Flight' comes from
The time of flight of each 1+ ion in this tube depends on their velocity
Again, this is important to remember when completing calculations

Stage 4: Detection
Once they have pass through the mass spectrometer, the 1+ ions will hit a negatively charged
'detector' plate
As they hit this electric plate, they gain an electron
This gaining of an electron discharges the ion, and causes a current to be produced

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This size of the current is proportional to the abundance of those ions hitting the plate and gaining
an electron
The detector plate is connected to a computer, which produces the mass spectrum Your notes

Key Equations for Time of Flight Mass Spectrometry

where
KE = kinetic energy of the particles (J)
m = mass of the particles (kg)
v = velocity of the particles (ms-1)
t = time of flight of the particles (s)
d = the length of the flight tube (m)

Examiner Tips and Tricks

Remember: All particles in the mass spectrometer are accelerated to the same kinetic energy.
The time of flight is proportional to the square root of the mass of the ions, showing that the lighter
the ion the faster it will pass through and the quicker it will hit the detector.
The heavier the ion, the slower it will travel and the longer it will take to hit the detector.

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Shells & Orbitals

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Electron Shells
Shells
The arrangement of electrons in an atom is called the electron configuration
Electrons are arranged around the nucleus in principal energy levels or principal quantum shells
Principal quantum numbers (n) are used to number the energy levels or quantum shells
The lower the principal quantum number, the closer the shell is to the nucleus
So, the first shell which is the closest to the nucleus is n = 1
The higher the principal quantum number, the greater the energy of the shell and the further away
from the nucleus
Each principal quantum number has a fixed number of electrons it can hold
n = 1 : up to 2 electrons
n = 2 : up to 8 electrons
n = 3 : up to 18 electrons
n = 4 : up to 32 electrons

Subshells
Small PDF
The principal quantum shells are split into subshells which are given the letters s, p and d
Elements with more than 57 electrons also have an f shell
The energy of the electrons in the subshells increases in the order s < p < d
The order of subshells appears to overlap for the higher principal quantum shells as seen in the diagram
below:

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Electrons are arranged in principal quantum shells, which are numbered by principal quantum numbers
Orbitals
Subshells contain one or more atomic orbitals
Orbitals exist at specific energy levels and electrons can only be found at these specific levels, not in
between them
Each atomic orbital can be occupied by a maximum of two electrons
This means that the number of orbitals in each subshell is as follows:

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s : one orbital (1 x 2 = total of 2 electrons)

p : three orbitals ( 3 x 2 = total of 6 electrons) Your notes
d : five orbitals (5 x 2 = total of 10 electrons)
f : seven orbitals (7 x 2 = total of 14 electrons)
The orbitals have specific 3-D shapes
s orbital shape
The s orbitals are spherical
The size of the s orbitals increases with increasing shell number
E.g. the s orbital of the third quantum shell (n = 3) is bigger than the s orbital of the first quantum
shell (n = 1)

p orbital shape
The p orbitals have a dumbbell shape
Every shell has three p orbitals except for the first one (n = 1)
The p orbitals occupy the x, y and z axes and point at right angles to each other, so are oriented
perpendicular to one another
The lobes of the p orbitals become larger and longer with increasing shell number

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Your notes

Representation of orbitals (the dot represents the nucleus of the atom) showing spherical s orbitals (a),
p orbitals containing ‘lobes’ along the x, y and z axis

Note that the shape of the d orbitals is not required

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Your notes

An overview of the shells, subshells and orbitals in an atom

Ground state
The ground state is the most stable electronic configuration of an atom which has the lowest amount
of energy
This is achieved by filling the subshells of energy with the lowest energy first (1s)
The order of the subshells in terms of increasing energy does not follow a regular pattern at n = 3 and
higher

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Your notes

The ground state of an atom is achieved by filling the lowest energy subshells first

Electron Arrangement Summary

Each shell can be divided further into subshells, labelled s, p, d and f
Each subshell can hold a specific number of orbitals:
s subshell : 1 orbital
p subshell : 3 orbitals
d subshell : 5 orbitals
f subshell : 7 orbitals
Each orbital can hold a maximum number of 2 electrons so the maximum number of electrons in each
subshell are as follows:
s : 1 x 2 = total of 2 electrons
p : 3 x 2 = total of 6 electrons
d : 5 x 2 = total of 10 electrons
f : 7 x 2 = total of 14 electrons

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Summary of the Arrangement of Electrons in Atoms Table

Main Energy Sub Number of orbitals Total number of Total number of Your notes
Level (n) Shells in sub-shell electrons in each orbital electrons in main shell

1 s 1 2 2

2 s 1 2 8

p 3 6

3 s 1 2 18

p 3 6

d 5 10

4 s 1 2 32

p 3 6

d 5 10

f 7 14

Examiner Tips and Tricks

The three p orbitals are labelled px, py and pz, but you do not need to include this in your electron
configurations!

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Electron Configuration
Your notes
Deducing the Electron Configuration
The Periodic Table is split up into four main blocks depending on their electron configuration
Elements can be classified as an s-block element, p-block element and so on, based on the position
of the outermost electron:
s block elements
Have their valence electron(s) in an s orbital
p block elements
Have their valence electron(s) in a p orbital
d block elements
Have their valence electron(s) in a d orbital
f block elements
Have their valence electron(s) in an f orbital

The principal quantum shells increase in energy with increasing principal quantum number

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E.g. n = 4 is higher in energy than n = 2

The subshells increase in energy as follows: s ＜ p ＜ d ＜ f
Your notes
The only exception to these rules is the 3d orbital which has slightly higher energy than the 4s
orbital
Because of this, the 4s orbital is filled before the 3d orbital
All the orbitals in the same subshell have the same energy and are said to be degenerate
E.g. px, py and pz are all equal in energy

Relative energies of the shells and subshells

The electron configuration gives information about the number of electrons in each shell, subshell and
orbital of an atom
The subshells are filled in order of increasing energy

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Your notes

The electron configuration shows the number of electrons occupying a subshell in a specific shell
Writing out the electron configuration tells us how the electrons in an atom or ion are arranged in their
shells, subshells and orbitals
This can be done using the full electron configuration or the shorthand version
The full electron configuration describes the arrangement of all electrons from the 1s subshell up
The shorthand electron configuration includes using the symbol of the nearest preceding noble
gas to account for however many electrons are in that noble gas
Ions are formed when atoms lose or gain electrons
Negative ions are formed by adding electrons to the outer subshell
Positive ions are formed by removing electrons from the outer subshell
The transition metals fill the 4s subshell before the 3d subshell but lose electrons from the 4s first
and not from the 3d subshell (the 4s subshell is lower in energy

Full Electron Configurations

Hydrogen has 1 single electron
The electron is in the s orbital of the first shell
Its electron configuration is 1s1
Potassium has 19 electrons
The first 2 electrons fill the s orbital of the first shell
They then continue to fill subsequent orbitals and subshells in order of increasing energy
The 4s orbital is lower in energy than the 3d subshell, so it is therefore filled first
The full electron configuration of potassium is 1s2 2s2 2p6 3s2 3p6 4s1

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Shorthand Electron Configurations

Using potassium as an example again: Your notes
The nearest preceding noble gas to potassium is argon
This accounts for 18 electrons of the 19 electrons that potassium has
The shorthand electron configuration of potassium is [Ar] 4s1

Worked Example
Write down the full and shorthand electron configuration of the following elements:
1. Calcium
2. Gallium
3. Mg2+
Answer 1:
Calcium has has 20 electrons so the full electronic configuration is:
1s2 2s2 2p6 3s2 3p6 4s2
The 4s orbital is lower in energy than the 3d subshell and is therefore filled first
The shorthand version is [Ar] 4s2 since argon is the nearest preceding noble gas to calcium
which accounts for 18 electrons
Answer 2:
Gallium has 31 electrons so the full electronic configuration is:
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p1
The shorthand electronic configuration is:
[Ar] 3d10 4s2 4p1
Even though the 4s is filled first, the full electron configuration is often written in numerical order.
So, if there are electrons in the 3d sub-shell, then these will be written before the 4s
Answer 3:
A magnesium atom has 12 electrons so its electronic configuration would be
1s2 2s2 2p6 3s2
To form a magnesium ion, it loses its two outer electrons so the electronic configuration for the
ion is:
1s2 2s2 2p6
Using the shorthand, the electronic configuration is:

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[Ne]

Exceptions Your notes

Chromium and copper have the following electron configurations, which are different to what you may
expect:
Cr is [Ar] 3d5 4s1 not [Ar] 3d4 4s2
Cu is [Ar] 3d10 4s1 not [Ar] 3d9 4s2
This is because the [Ar] 3d5 4s1 and [Ar] 3d10 4s1 configurations are energetically stable

Presenting the Electron Configuration

Electrons can be imagined as small spinning charges which rotate around their own axis in either a
clockwise or anticlockwise direction
The spin of the electron is represented by its direction
Electrons with similar spin repel each other which is also called spin-pair repulsion
Electrons will therefore occupy separate orbitals in the same subshell where possible, to minimize this
repulsion and have their spin in the same direction
E.g. if there are three electrons in a p subshell, one electron will go into each px, py and pz orbital

Electron configuration: three electrons in a p subshell

Electrons are only paired when there are no more empty orbitals available within a subshell, in which
case the spins are the opposite spins to minimize repulsion
E.g. if there are four electrons in a p subshell, one p orbital contains 2 electrons with opposite spin
and two orbitals contain one electron only
The first 3 electrons fill up the empty p orbitals one at a time, and then the 4th one pairs up in the px
orbital

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Your notes

Electron configuration: four electrons in a p subshell

Box Notation
The electron configuration can be represented using the electrons in boxes notation
Each box represents an atomic orbital
The boxes are arranged in order of increasing energy from bottom to top
The electrons are represented by opposite arrows to show the spin of the electrons
E.g. the box notation for titanium is shown below
Note that since the 3d subshell cannot be either full or half full, the second 4s electron is not
promoted to the 3d level and stays in the 4s orbital

The electrons in titanium are arranged in their orbitals as shown. Electrons occupy the lowest energy
levels first before filling those with higher energy

Examiner Tips and Tricks

You can use full headed arrows or half headed arrows to represent electrons in your box notations.

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Ionisation Energy
Your notes
What is Ionisation Energy?
The Ionisation Energy (IE) of an element is the amount of energy required to remove one mole of
electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions
Ionisation energies are measured under standard conditions which are 298 K and 101 kPa
The units of IE are kilojoules per mole (kJ mol-1)
The first ionisation energy (IE1) is the energy required to remove one mole of electrons from one mole
of atoms of an element to form one mole of 1+ ions
E.g. the first ionisation energy of gaseous calcium:
Ca(g) → Ca+ (g) + e- IE1 = +590 kJ mol-1

Trends in Ionisation Energies

Ionisation energies show periodicity - a trend across a period of the Periodic Table
As could be expected from their electron configuration, the group 1 metals have a relatively low
ionisation energy, whereas the noble gases have very high ionisation energies
The size of the first ionisation energy is affected by four factors:
Size of the nuclear charge
Distance of outer electrons from the nucleus
Shielding effect of inner electrons
Spin-pair repulsion
First ionisation energy increases across a period and decreases down a group

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Your notes

A graph showing the ionisation energies of the elements hydrogen to sodium

Ionisation energy across a period
The ionisation energy across a period generally increases due to the following factors:
Across a period the nuclear charge increases
This causes the atomic radius of the atoms to decrease, as the outer shell is pulled closer to the
nucleus, so the distance between the nucleus and the outer electrons decreases
The shielding by inner shell electrons remain reasonably constant as electrons are being added to
the same shell
It becomes harder to remove an electron as you move across a period; more energy is needed
So, the ionisation energy increases

Dips in the trend

There is a slight decrease in IE1 between beryllium and boron as the fifth electron in boron is in the 2p
subshell, which is further away from the nucleus than the 2s subshell of beryllium
Beryllium has a first ionisation energy of 900 kJ mol-1 as its electron configuration is 1s2 2s2
Boron has a first ionisation energy of 800 kJ mol-1 as its electron configuration is 1s2 2s2 2px1

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There is a slight decrease in IE1 between nitrogen and oxygen due to spin-pair repulsion in the 2px
orbital of oxygen
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Nitrogen has a first ionisation energy of 1400 kJ mol-1 as its electron configuration is 1s2 2s2 2px1
2py1 2pz1
Oxygen has a first ionisation energy of 1310 kJ mol-1 as its electron configuration is 1s2 2s2 2px2 2py1
2pz1
In oxygen, there are 2 electrons in the 2px orbital, so the repulsion between those electrons makes
it slightly easier for one of those electrons to be removed
From one period to the next
There is a large decrease in ionisation energy between the last element in one period, and the first
element in the next period
This is because:
There is increased distance between the nucleus and the outer electrons as you have added a new
shell
There is increased shielding by inner electrons because of the added shell
These two factors outweigh the increased nuclear charge

Ionisation energy down a group

The ionisation energy down a group decreases due to the following factors:
The number of protons in the atom is increased, so the nuclear charge increases
But, the atomic radius of the atoms increases as you are adding more shells of electrons, making
the atoms bigger
So, the distance between the nucleus and outer electron increases as you descend the group
The shielding by inner shell electrons increases as there are more shells of electrons
These factors outweigh the increased nuclear charge, meaning it becomes easier to remove the
outer electron as you descend a group
So, the ionisation energy decreases

Ionisation Energy Trends across a Period & going down a Group Table
Across a period: Ionisation energy increases Down a group: Ionisation energy decreases

Increase in nuclear charge Increase in nuclear charge

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Same number of shells More shells

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Distance from outer electron to the nucleus Distance from outer electron to the nucleus
decreases increases

Shielding remains reasonably constant Shielding increases, cause a weaker force of

attraction between the outer electron and the
nucleus

Decreased atomic / ionic radius Increased atomic / ionic radius

The outer electron is held more tightly to the The outer electron is held less tightly to the nucleus
nucleus so it requires more energy to remove so it requires less energy to remove

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Ionisation Energy: Trends & Evidence

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Ionisation Energies: Equations
The second ionisation energy (IE2) is the energy required to remove the second mole of electrons from
each +1 ion in a mole of gaseous +1 ions, to form one mole of +2 ions
The third ionisation energy (IE3) is the energy required to remove the third mole of electrons from each
+2 ion in a mole of gaseous +2 ions, to form one mole of +3 ions
And so on...
The electrons from an atom can be continued to be removed until only the nucleus is left
This sequence of ionisation energies is called successive ionisation energies

Successive Ionisation Energies of Beryllium Table

Ionisation energy Equation

First Be (g) → Be+ (g) + e-

Second Be+ (g) → Be2+ (g) + e-

Third Be2+ (g) → Be3+ (g) + e-

Fourth Be3+ (g) → Be4+ (g) + e-

Examiner Tips and Tricks

Remember: Equations representing ionisation energies must have gaseous (g) state symbols for
the atoms and ions but not for the electrons.
You will lose the mark in your exam if you do not include the state symbols, even if the question does
not specify for you to include them.

Successive Ionisation Energies

Successive ionisation energies of an element
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The successive ionisation energies of an element increase

This is because once you have removed the outer electron from an atom, you have formed a positive Your notes
ion
Removing an electron from a positive ion is more difficult than from a neutral atom
As more electrons are removed, the attractive forces increase due to decreasing shielding and an
increase in the proton to electron ratio
The increase in ionisation energy, however, is not constant and is dependent on the atom's electronic
configuration
Taking calcium as an example:

Ionisation Energies of Calcium Table

Electronic configuration 1s2 2s2 2p6 1s2 2s2 2p6 1s2 2s2 2p6 1s2 2s2 2p6
3s2 3p6 4s2 3s2 3p6 4s1 3s2 3p6 3s2 3p5

Ionisation energy First Second Third Fourth

Ionisation energy 590 1150 4940 6480

(kJ mol-1)

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The first electron removed has a low IE1 as it is easily removed from the atom due to the spin-pair
repulsion of the electrons in the 4s orbital
The second electron is more difficult to remove than the first electron as there is no spin-pair repulsion
The third electron is much more difficult to remove than the second one corresponding to the fact that
the third electron is in a principal quantum shell which is closer to the nucleus (3p)
Removal of the fourth electron is more difficult as the orbital is no longer full, and there is less spin-pair
repulsion
The graph shows there is a large increase in successive ionisation energy as the electrons are being
removed from an increasingly positive ion
The big jumps on the graph show the change of shell and the small jumps are the change of subshell

Examiner Tips and Tricks

It gets more difficult to remove electrons from principal quantum shells that get closer to the
nucleus, as there is less shielding and an increase in attractive forces between the electrons and
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nuclear charge.
Be careful with interpreting successive ionisation energy graphs, especially if you are not given Your notes
every successive ionisation energy and are just shown part of the graph - you should count the
electrons from left to right!
It is a good idea to label the shells and subshells on ionisation energy graphs in an exam so that you
do not make the mistake of reading the graph backwards.

Successive ionisation data can be used to:

Predict or confirm the simple electronic configuration of elements
Confirm the number of electrons in the outer shell of an element
Deduce the Group an element belongs to in the Periodic Table
By analyzing where the large jumps appear and the number of electrons removed when these large
jumps occur, the electron configuration of an atom can be determined
Na, Mg and Al will be used as examples to deduce the electronic configuration and positions of
elements in the Periodic Table using their successive ionisation energies

Successive Ionisation Energies Table

Element Atomic number First ionisation energy (kJ mol-1)

First Second Third Fourth

Na 11 494 4560 6940 9540

Mg 12 736 1450 7740 10500

Al 13 577 1820 2740 11600

Sodium
For sodium, there is a huge jump from the first to the second ionisation energy, indicating that it is
much easier to remove the first electron than the second
Therefore, the first electron to be removed must be the last electron in the valence shell thus Na
belongs to group I
The large jump corresponds to moving from the 3s to the full 2p subshell
Na 1s2 2s2 2p6 3s1
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Magnesium
There is a huge increase from the second to the third ionisation energy, indicating that it is far easier to Your notes
remove the first two electrons than the third
Therefore the valence shell must contain only two electrons indicating that magnesium belongs to
group II
The large jump corresponds to moving from the 3s to the full 2p subshell
Mg 1s2 2s2 2p6 3s2

Aluminium
There is a huge increase from the third to the fourth ionisation energy, indicating that it is far easier to
remove the first three electrons than the fourth
The 3p electron and 3s electrons are relatively easy to remove compared with the 2p electrons which
are located closer to the nucleus and experience greater nuclear charge
The large jump corresponds to moving from the third shell to the second shell
Al 1s2 2s2 2p6 3s2 3p1

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