Physical properties of the Group 2 elements

Some of the most useful physical properties of the Group 2 metals are shown in below

Element Atomic radius/nm 1st ionisation energy/kJ mol−1 Electronegativity Melting point/°C
Beryllium 0.111 900 1.57 1278
Magnesium 0.160 738 1.31 649
Calcium 0.197 590 1.00 839
Strontium 0.215 550 0.95 769
Barium 0.217 503 0.89 729
  • The first three physical properties show steady trends – increasing upwards in atomic radius and downwards in first ionisation energy and electronegativity. The decrease in melting point would fit this pattern if it were not for the anomalous low value for magnesium.
  •  You can explain the change in atomic radius in terms of the additional shell of electrons added for each period and the reduced effective nuclear charge as more electron shells are added.
  • As Group 2 is descended, the increase in charge of the nucleus is offset by the number of inner electrons. However, the distance of the outer electrons from the nucleus increases and the first ionisation energy decreases down the group.
  • The electronegativity of the atoms decreases down the group. As the size of the atoms increases, any bonding pair of electrons is further from the nucleus, which means it is less strongly held and the electronegativity decreases. The effect of this is to increase the ionic (electrovalent) character of any compounds as the group is descended.

Reactions of the Group 2 elements with oxygen, water and dilute acids

  • The only reactions you need to remember for the Group 2 elements are their reactions with oxygen, water, and dilute hydrochloric and sulfuric acids
Element Reaction with oxygen Reaction with water Reaction with dilute acids
Beryllium Reluctant to burn, white flame No reaction Reacts rapidly
Magnesium Burns easily with a bright white flame Reacts vigorously with steam but only slowly with water Reacts vigorously
Calcium Difficult to ignite, flame tinged red Reacts moderately forming the hydroxide Reacts vigorously
Strontium Difficult to ignite, flame tinged red Reacts rapidly forming the hydroxide Reacts violently
Barium Difficult to ignite, flame tinged green Reacts vigorously forming the hydroxide Reacts violently
  • energy and electronegativity. The decrease in melting point would fit this pattern if it were not for the anomalous low value for magnesium.
  • You can explain the change in atomic radius in terms of the additional shell of electrons added for each period and the reduced effective nuclear charge as more electron shells are added.
  • As Group 2 is descended, the increase in charge of the nucleus is offset by the number of inner electrons. However, the distance of the outer electrons from the nucleus increases and the first ionisation energy decreases down the group.
  • The electronegativity of the atoms decreases down the group. As the size of the atoms increases, any bonding pair of electrons is further from the nucleus, which means it is less strongly held and the electronegativity decreases. The effect of this is to increase the ionic (electrovalent) character of any compounds as the group is descended.

Reactions of the Group 2 elements with oxygen, water and dilute acids

The only reactions you need to remember for the Group 2 elements are their reactions with oxygen, water, and dilute hydrochloric and sulfuric acids.

Element Reaction with oxygen Reaction with water Reaction with dilute acids
Beryllium Reluctant to burn, white flame No reaction Reacts rapidly
Magnesium Burns easily with a bright white flame Reacts vigorously with steam but only slowly with water Reacts vigorously
Calcium Difficult to ignite, flame tinged red Reacts moderately forming the hydroxide Reacts vigorously
Strontium Difficult to ignite, flame tinged red Reacts rapidly forming the hydroxide Reacts violently
Barium Difficult to ignite, flame tinged green Reacts vigorously forming the hydroxide Reacts violently

The general equation for the reaction with oxygen is:

where X is any metal in the group.

Both strontium and barium can also form a peroxide as well as the oxide:

where Y is strontium or barium.

The general equation for the reaction with water is:

The exception to this is magnesium, which forms the oxide when reacted with steam.

The general equation for the reaction with dilute acids is:

Behaviour of Group 2 oxides, hydroxides and carbonates with water and dilute acids

Beryllium oxide is amphoteric, but all the other oxides are sparingly soluble in water, producing solutions of increasing base strength:

The hydroxides increase in solubility down the group, due to the decrease in lattice dissociation enthalpy, which outweighs the change in the enthalpy of hydration of the metal ion.

The carbonates decrease in solubility down the group, due to the decrease in the enthalpy of hydration of the metal ion.

The compounds react with dilute hydrochloric and sulfuric acids depending on the solubility of the salts they produce. Only the magnesium compounds react appreciably with sulfuric acid because the other sulfates are sparingly soluble.

Thermal decomposition of Group 2 nitrates and carbonates

The changes in thermal stability stem from of the ability of a cation to polarise the anion. This is more pronounced at the top of the group, where the cations are smaller and have a high charge density. This applies to both the nitrate and carbonate, where polarisation results in the formation of the oxide:

2X(NO3)2(s) → 2XO(s) + 4NO2(g) + O2(g)

XCO3(s) → XO(s) + CO2(g)

Solubility of Group 2 sulfates and hydroxides

The solubility and the enthalpy change of solution of the sulfates of Group 2 elements decrease down the group. This is due to a combination of the relative sizes of the enthalpy change of hydration of the cations and the lattice energy of the sulfate concerned:

The solubility of the sulfates of Group 2 decreases down the group. This is due to the relative magnitudes of the enthalpy change of hydration and the lattice energy for compounds.

  • As the cations get bigger, the energy released when the ions bond to water molecules (the enthalpy change of hydration) falls.
  • Larger ions are not as strongly attracted to the water molecules.
  • As you go down a group, the energy needed to break up the lattice decreases as the positive ions get bigger. The bigger the ions, the more distance there is between them and the weaker are the forces holding them together.
  • Because both energy changes decrease, it is a question of which is the more significant. For large ions, such as SO42−, it is the enthalpy change of hydration factor that dominates.
  • Conversely the hydroxides of Group 2 elements become more soluble descending the group, but there is not a simple explanation for this.

Practice Questions

  1. Describe and explain qualitatively the trend in thermal stability of Group 2 nitrates and carbonates using the effect of ionic radius on the polarisation of a large anion
  2. Describe and explain qualitatively the variation in solubility and of enthalpy change of solution, ΔHsol, of Group 2 hydroxides and sulfates in terms of the magnitudes of the enthalpy change of hydration and the lattice energy.
  3. Predict the relative thermal stabilities of MgCO3 and BaCO3, giving a reason for your answer.

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