Topic 8 · page 9 of 29

Ionisation energy and reactivity down Groups 1 and 2

Down both groups the outer electron is further from the nucleus and more shielded, so ionisation energy falls. Group 1 and 2 reactions are oxidations of the metal, so the metals become more reactive down the group.

13 min read8.118.12

Read first: Oxidising and reducing agents

First ionisation energy 8.11

The second ionisation energy is the energy to remove a second electron, from each ion in one mole of gaseous 1+ ions: MX+(g)→MX2+(g)+eX−\ce{M+(g) -> M^2+(g) + e-}. It is always larger than the first, because the electron is removed from an ion that is already positive.

Group 1 Li Na K Rb Cs
First ionisation energy / kJ mol⁻¹ 520 496 419 403 376
Group 2 Be Mg Ca Sr Ba
First / kJ mol⁻¹ 900 738 590 550 503
Second / kJ mol⁻¹ 1757 1451 1145 1064 965
First + second / kJ mol⁻¹ 2657 2189 1735 1614 1468
Two bar charts. Left, Group 1 first ionisation energies in kilojoules per mole: lithium 520, sodium 496, potassium 419, rubidium 403, caesium 376, the bars getting shorter down the group. Right, Group 2 first and second ionisation energies: beryllium 900 and 1757, magnesium 738 and 1451, calcium 590 and 1145, strontium 550 and 1064, barium 503 and 965, both falling down the group, with the second always larger than the first.
Ionisation energy falls down both groups. Data-book values, rounded.

The trend and its explanation

Down each group the first ionisation energy decreases. Three things change down a group, and one effect wins.

  1. The outer electron is in a higher shell, so it is further from the nucleus. The attraction falls with distance.
  2. There are more inner shells between the outer electron and the nucleus. These shield it: they repel the outer electron and cut down the pull of the nucleus.
  3. The nuclear charge also increases, which pulls harder. But distance and shielding together outweigh it.

Group 2 loses two electrons 8.11

A Group 2 atom forms M²⁺, so both ionisation energies matter. Compare the sum of the first and second, in the last row of the table. It also falls steadily: from 2189 kJ mol⁻¹ for magnesium to 1468 kJ mol⁻¹ for barium. The explanation is the same: each outer electron is further out and more shielded.

Across a period the pattern is reversed. A Group 2 element has a higher first ionisation energy than the Group 1 element beside it (Mg 738, Na 496), because it has a higher nuclear charge acting on electrons in the same shell.

Why reactivity increases down the group 8.12

Every reaction of a Group 1 or 2 metal is an oxidation of the metal:

M→MX++eX−M→MX2++2 eX− \ce{M -> M+ + e-} \qquad\qquad \ce{M -> M^2+ + 2e-}

The metal is a reducing agent, and the more easily it gives away its outer electrons, the more reactive it is. So the argument has four links:

  1. Down the group the outer electrons are further away and more shielded.
  2. They are held less strongly, so less energy is needed to remove them: ionisation energy falls.
  3. The metal loses electrons more easily: it is oxidised more readily.
  4. So the metal is more reactive: Li < Na < K and Mg < Ca < Sr < Ba.

Worked exampleAn exam answer

Explain why potassium is more reactive than sodium.

  1. Both lose one electron to form a 1+ ion (0 → +1): the reaction is an oxidation.
  2. The outer electron in potassium is in a higher shell, further from the nucleus, and is more shielded by the extra inner shell.
  3. The attraction between the nucleus and the outer electron is weaker, so the first ionisation energy of potassium (419 kJ mol⁻¹) is lower than that of sodium (496 kJ mol⁻¹).
  4. Potassium loses its electron more easily, so it reacts more readily.

Check yourself

  1. Define first ionisation energy, and write an equation with state symbols for the first ionisation of sodium.
  2. Explain why the first ionisation energy of calcium is lower than that of magnesium.
  3. Why does the nuclear charge not make ionisation energy rise down the group?
  4. The first ionisation energy of barium is 503 kJ mol⁻¹ and of strontium 550 kJ mol⁻¹. State which is more reactive and explain.
  5. Calculate the energy needed to form 1 mol of Mg²⁺(g) from 1 mol of Mg(g), using 738 and 1451 kJ mol⁻¹. Why is this the relevant figure for reactions of magnesium?
  6. A student writes "potassium is more reactive than lithium because it has more electrons to lose". Correct the statement.

Answers cover these until you have tried the questions

  1. The energy needed to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions. Na(g)→NaX+(g)+eX−\ce{Na(g) -> Na+(g) + e-}
  2. Calcium's outer electrons are in a higher shell (further from the nucleus) with more inner shells shielding them. The attraction is weaker, so less energy is needed to remove one.
  3. The nuclear charge does increase, but the extra shells increase the distance and the shielding. Together these outweigh the extra nuclear charge, so the net attraction on the outer electron is weaker.
  4. Barium is more reactive. Its outer electrons are further from the nucleus and more shielded, so its ionisation energy (503) is lower than strontium's (550). It is oxidised more easily.
  5. 738 + 1451 = 2189 kJ mol⁻¹. Magnesium forms Mg²⁺ in its reactions (Mg: 0 → +2), so both electrons are removed.
  6. Both lithium and potassium lose one electron. Potassium is more reactive because its outer electron is further away and more shielded, so it needs less energy to remove and is oxidised more easily.