Trends in Ionisation Energy Chemistry Tutorial

Key Concepts

• Ionization energy is also referred to as ionisation energy.
• Ionization energy is given a number of symbols including I, E, and I.E.
• Ionization energy is measured in kilojoules per mole (kJ mol^-1) or electronvolts per atom (eV)¹.
• Ionization energy, or ionisation energy, is the energy required to remove an electron from a gaseous atom or ion.
• First Ionization energy is the energy required to remove an electron from the gaseous atom

  First Ionization for the element M:

  M_(g) → M⁺_(g) + e^-, first ionization energy is I₁

  First Ionization for Hydrogen:

  H_(g) → H⁺_(g) + e^-

  First Ionization for Carbon:

  C_(g) → C⁺_(g) + e^-

• Second Ionization energy is the energy required to remove an electron from the gaseous ion

  Second Ionization of element M:

  M⁺_(g) → M²⁺_(g) + e^-, second ionization energy is I₂

  Second Ionization for Carbon:

  C⁺_(g) → C²⁺_(g) + e^-

• The second ionization energy of an element will be higher than the first ionization energy.

  For an element: I₂ > I₁

• General trends in the ionization energy of elements in the Periodic Table:

  ⚛ Ionisation energy decreases down a group.

  ⚛ Ionisation energy increases across a period from left to right.

• Successive ionization energies for an element provide evidence for the number of electrons occupying the highest energy level, or valence shell, of the atom.

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Trends in First Ionization Energy (Ionisation Energy) in Groups of the Periodic Table

As you go down a Group in the Periodic Table from top to bottom, the electron being removed occupies a higher energy level and is therefore further away from the nucleus.
The force of attraction between the negatively charged electron being removed and the positively charged nucleus decreases as you go down the Group.

Trends in the First Ionisation Energy of Group 1 (IA, Alkali Metals) Elements

Consider the first ionisation energy of the elements in Group 1 of the Periodic Table as shown in the table below:

First ionization energy decreases as you go down Group 1 because the electron being removed is further from the nucleus (in a higher energy level).

Trends in the First Ionisation Energy of Group 17 (VIIA, Halogens) Elements

Consider the first ionisation energy of the elements in Group 17 of the Periodic Table as shown in the table below:

First ionization energy decreases as you go down Group 17 because the electron being removed is further from the nucleus (in a higher energy level).

Trends in First Ionisation Energy of Group 18 (VIIIA, 0, Noble Gases) Elements

Consider the first ionisation energy of the elements in Group 18 of the Periodic Table as shown in the table below:

First ionization energy decreases as you go down Group 18 because the electron being removed is further from the nucleus (in a higher energy level).

Trends in First Ionization Energy in Periods of the Periodic Table

In general, the first ionisation energy of elements increases as you go across a Period from left to right.
As the nuclear charge increases across a period, the electron being removed is more strongly bound to the nucleus, and, as the atomic radius decreases, the negatively charged electron being removed is closer to the positively charged nucleus.

Irregularities in this trend are due to the location of the electron being removed in terms of its orbital and how many electrons are present in that orbital.

Trends in the First Ionisation Energy Across Period 2

Consider the first ionisation energy of the elements across Period 2 of the Periodic Table as shown in the table below:

Ionization energy generally increases across Period 2 from left to right as the increasing nuclear charge on successive atoms more tightly binds the electron being removed to the nucleus.

Trends in the First Ionisation Energy Across Period 3

Consider the first ionisation energy of the elements across Period 3 of the Periodic Table as shown in the table below:

Ionization energy generally increases across Period 3 from left to right as the increasing nuclear charge on successive atoms more tightly binds the electron being removed to the nucleus.

Trends in Successive Ionization Energy of Elements in the Periodic Table

• Second ionisation energy is greater than first ionisation energy for an element as it is harder to remove a negatively charged electron from the positively charged ion.
• Third ionization energy is greater than second ionization energy for an element because you are now trying to remove a negatively charged electron from an ion with a 2+ charge.
• Removing successive electrons from the same energy level requires a little more energy each time, but removing an electron from a lower energy level requires much, much, more energy because the electron being removed is then so much closer to the positively charged nucleus.

Successive Ionisation Energies for Group 1 Elements (IA, Alkali Metals)

Consider the first, second and third ionisation energies of Group 1 elements as shown in the table below:

• Going down the group, the trend is that ionization energy decreases because as you go down the group you are removing an electron from a higher energy level which is further from the nucleus.

  First ionization energy decreases as you go down the group.

  Second ionization energy decreases as you go down the group.

  Third ionization energy decreases as you go down the group.

• For each element in the Group, the first ionization energy is less than the second ionization energy which is less than the third ionization energy.

  I₁ < I₂ < I₃

  Each time an electron is removed, it results in there being 1 more proton in the nucleus with its positive charge not being balanced by a negatively charged electron, so the relative nuclear charge is increasing.

  As this relative nuclear charge increases, the remaining electrons are more strongly attracted to the nucleus making it even harder to remove the next electron.

• The ratio of the first ionization energy to the second is much, much, less than the ratio of the second ionization energy to the third.

  I₁ << I₂ < I₃
  or
  I₁/I₂ << I₂/I₃

  It is much, much, easier to remove one electron from a Group 1 atom than it is to remove an electron from a Group 1 ion with charge +1.

  This suggests that the second electron being removed is in a lower energy level and therefore closer to the nucleus and much more strongly attracted to it.

  The first electron being removed must be in a higher energy level.

Successive Ionisation Energies for Period 2 Elements

Consider the successive ionisation energies of Period 2 elements as shown in the table below:

• For each element: I₁ < I₂ < I₃ etc
  as it becomes increasingly difficult to remove a negatively charged electron from a positively charged ion of charge +1 then +2 etc
• Removing electrons from a higher energy level is much easier than removing electrons from the lower energy levels which are closer to the nucleus:

  Li(g)	→	Li+(g) + e-		I1 = 520 kJ mol-1
  1s22s1	→	1s2
  Li+(g)	→	Li2+(g) + e-		I2 = 7300 kJ mol-1	I2/I1 = 14
  1s2	→	1s1
  Removing electrons from the first energy level (1s electrons) is much harder than removing electrons from the second energy level (2s electrons).
  Li2+(g)	→	Li3+(g) + e-		I3 = 10950 k mol-1	I3/I2 = 1.5
  1s1	→	Li nucleus
  Removing the second 1s electron is only slightly harder than removing the first.

  Be(g)	→	Be+(g) + e-		I1 = 899 kJ mol-1
  1s22s2	→	1s22s1
  Be+(g)	→	Be2+(g) + e-		I2 = 1721 kJ mol-1	I2/I1 = 1.9
  1s22s1	→	1s2
  I2 is only slightly greater than I1 because both electrons are being removed from the same energy level.
  Be2+(g)	→	Be3+(g) + e-		I3 = 14513 kJ mol-1	I3/I2 = 8.4
  1s2	→	1s1
  Removing an electron from the first energy level is much, much harder than removing the electrons from the second energy level.
  Be3+(g)	→	Be4+(g) + e-		I4 = 20550 kJ mol-1	I3/I2 = 1.4
  1s1	→	Be nucleus
  Removing the second electron from the first energy is only slighly harder than removing the first one.