AQA AS Chemistry

Revision Notes

1.1.7 Ionisation Energy: Trends & Evidence

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

Atomic Structure Ionisation Energy across Periods, downloadable AS & A Level Chemistry revision 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
  • There is a slight decrease in IE1 between nitrogen and oxygen due to spin-pair repulsion in the 2px orbital of oxygen
    • 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

Atomic Structure First Ionisation Energy Trends Table, downloadable AS & A Level Chemistry revision notes

Successive ionisation energies of an element

  • 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 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

Atomic Structure First Four Ionisation Energies of Calcium Table, downloadable AS & A Level Chemistry revision notes

Ionization Energies_ Trends Successive Ionisation Energies of Calcium_1, downloadable IB Chemistry revision notes

  • 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

Exam Tip

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 nuclear charge.

Be careful with how you interpret successive ionisation energy graphs, especially if you are not given 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.

Ionisation Energy Trends: Explained

  • Energy is required to remove an outer shell electron as this involves breaking the attractive forces between the electron and the positively charged nucleus
  • There are several factors which affect the magnitude of the ionisation energy:
  • Nuclear charge
    • Positive nuclear charge increases with increasing number of protons
    • The greater the positive charge, the greater the attractive forces between the outer electron(s) and the nucleus
    • More energy is required to overcome these forces so ionisation energy increases with increasing nuclear charge
  • Shielding
    • Electrons repel each other and electrons occupying the inner shells repel electrons located in shells further outside the nucleus and prevent them from feeling the full effect of the nuclear charge
    • The greater the shielding effect is, the weaker the attractive forces between the positive nucleus and the negatively charged electrons
    • Less energy is required to overcome the weakened attractive forces so ionisation energy decreases with increasing shielding effects

Atomic Structure Ionisation & Shielding, downloadable AS & A Level Chemistry revision notes

Shielding makes it easier to remove the outermost electrons

  • Atomic/ionic radius
    • The larger the radius, the greater the distance between the nucleus and the outer shell electron(s)
    • Increasing distance weakens the strength of the attractive forces
    • Larger atoms/ions also result in greater shielding due to the presence of more inner electrons
    • Less energy is required to remove the outer shell electron(s) so ionisation energy decreases with increasing atomic/ionic radius
  • Spin-pair repulsion
    • Spin pair repulsion occurs when the electron being removed is spin paired with another electron in the same orbital
    • The proximity of the like charges of electrons in the orbital results in repulsion
    • Less energy is required to remove one of the electrons so ionisation energy decreases when there is spin-pair repulsion

Atomic radius

  • The atomic radius of an element is a measure of the size of an atom
  • Atomic radii show predictable patterns across the Periodic Table
    • They generally decrease across each Period
    • They generally increase down each Group
  • These trends can be explained by the electron shell theory
    • Atomic radii decrease as you move across a Period as the atomic number increases (increased positive nuclear charge) but at the same time extra electrons are added to the same principal quantum shell
    • The larger the nuclear charge, the greater the pull of the nuclei on the electrons which results in smaller atoms
    • Atomic radii increase moving down a Group as there is an increased number of shells going down the Group
    • The electrons in the inner shells repel the electrons in the outermost shells, shielding them from the positive nuclear charge
    • This weakens the pull of the nuclei on the electrons resulting in larger atoms

Atomic Structure Atomic Radius Trends, downloadable AS & A Level Chemistry revision notes

Trends in the atomic radii across a period and down a group

  • The diagram shows that the atomic radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period
  • This is because the alkali metals at the beginning of the next period have one extra principal quantum shell
    • This increases shielding of the outermost electrons and therefore increases the atomic radius

Ionic radius

  • The ionic radius of an element is a measure of the size of an ion
  • Ionic radii show predictable patterns
    • Ionic radii increase with increasing negative charge
    • Ionic radii decrease with increasing positive charge
  • These trends can also be explained by the electron shell theory
    • Ions with negative charges are formed by atoms accepting extra electrons while the nuclear charge remains the same
    • The outermost electrons are further away from the positively charged nucleus and are therefore held only weakly to the nucleus which increases the ionic radius
    • The greater the negative charge, the larger the ionic radius
    • Positively charged ions are formed by atoms losing electrons
    • The nuclear charge remains the same but there are now fewer electrons which undergo a greater electrostatic force of attraction to the nucleus which decreases the ionic radius
    • The greater the positive charger, the smaller the ionic radius

Atomic Structure Factors Affecting Ionisation Energies, downloadable AS & A Level Chemistry revision notes

Summary of factors affecting ionisation energies of atoms

Successive Ionisation Energies

  • 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

Atomic Structure First Four Ionisation Energies of Sodium, Magnesium & Aluminium Table, downloadable AS & A Level Chemistry revision notes

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

Magnesium

  • There is a huge increase from the second to the third ionisation energy, indicating that it is far easier to 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
  • This is due to weakened shielding effects through the loss of three electrons
  • The large jump corresponds to moving from the third shell to the second shell
    Al         1s2 2s2 2p6 3s2 3p1

Exam Tip

Find the large jumps by subtracting the successive ionisation energies from each other to identify when an electron has been removed from a different shell.

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