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 gaseous atoms of an element to form one mole of gaseous 1+ ions
  • E.g. the first ionisation energy of gaseous calcium:

   

Ca (g) → Ca+?(g) + e-? ? ? ? ??IE1?= +590 kJ mol-1

Successive ionisation energies of an element

• More than one electron can be removed form an atom and each time you remove an electron there is a successive ionisation energy
• These are called second, third ionisation energy and so on
• The?second ionisation energy?(IE2) is defined as
  • is the energy required to remove?one mole of electrons?from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions
• And can be represented as can be represented as

X+?(g) → X2+?(g) + e-

OR (for calcium)

Ca+?(g) → Ca2+?(g) + e-??IE2?= 1145 kJ mol-1

• The third ionisation energy?(IE3)?of?an element can be represented as

X2+?(g) → X3+?(g) + e-

OR (for calcium)

Ca2+?(g) → Ca3+?(g) + e-??IE3?= 4912 kJ mol-1

• 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

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

   

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