When we consider the elements of Group 2 we are invariably only investigating magnesium, calcium, strontium and barium. All isotopes of radium are radioactive and beryllium shows some strange properties for a metal as we will see. Each element has two outer shell / valence electrons.
| Be | [He] 2s2 |
| Mg | [Ne] 3s2 |
| Ca | [Ar] 4s2 |
| Sr | [Kr] 5s2 |
| Ba | [Xe] 6s2 |
As elements Group 2 metals have higher melting and boiling points than their Group 1 counterparts which is the result of stronger metallic bonding in the metal lattices. They have relatively low 1st and 2nd ionisation energies and so lose two valence electrons per atom to the the sea of delocalised electrons in the structure.
Mg(g) ⇾ Mg+(g) + e– 1st I.E. 738 kJ mol-1
Mg+(g) ⇾ Mg2+(g) + e– 2nd I.E. 1451 kJ mol-1
Group 2 metals form compounds in which they exist as M2+ cations and hence have an oxidation state of +2.
The strange case of beryllium
Beryllium is often seen to form covalent compounds which is highly unusual for a metal. This is because the Be2+ ion has a very high charge density and polarises the electron density of the anions it is forming a compound with. This gives the bonds a distinctly covalent character. You can find out more about ionic lattices and the grey areas between ionic, covalent and metallic bonding in this post.
In the case of beryllium carbonate the compound is innately unstable and decomposes to form beryllium oxide and carbon dioxide if not stored under carbon dioxide.
In beryllium chloride, BeCl2, the structure is similar to that of a long chain polymer with a mixture of single covalent and dative covalent bonds between the beryllium and chlorine atoms.
Explaining the reactivity of the Group 2 metals
Group 2 metals are much less reactive than Group 1 metals (only barium is stored under oil) and the oxide layer coating beryllium, magnesium, calcium and strontium makes them less reactive still. They all readily react with acids to give the metal salt and hydrogen.
Mg(s) + 2HCl(aq) ⇾ MgCl2(aq) + H2(g)
Beryllium is amphoteric (it can act as either an acid or a base) and in addition to reacting with acids, it will dissolve in a strong alkali to give the tetrahydroxoberyllate ion, [Be(OH)4]2-.
Be(s) + 4OH–(aq) ⇾ [Be(OH)4]2-(aq)
The reactivity of Group 2 metals increases as you descend the group (in simple terms, it takes less energy for barium to lose its valence electrons to form the Ba2+ ion as the 6s2 electrons are a long way from the positive nuclear charge and these valence electrons experience a weak electrostatic attraction). We can see this in practice as beryllium does not react with water, magnesium only reacts with water in the form of steam but calcium reacts vigorously with cold water.
Ca(s) + 2H2O(l) ⇾ Ca(OH)2(aq) + H2(g)
Compounds of Group 2
1. Group 2 oxides
Group 2 oxides are formed when the elements burn in air and the reaction becomes more exothermic as you descend the group – barium has been known to ignite simply in moist air!
2Sr(s) + O2(g) ⇾ 2SrO(s)
The oxides have high melting points (strong ionic bonds in the lattice) and they behave as bases, all bar beryllium oxide dissolving in water to form alkaline metal hydroxide solutions.
CaO(s) + H2O(l) ⇾ Ca(OH)2(aq)
Barium oxide will react further with excess oxygen at high temperatures to form barium peroxide.
2BaO(s) + O2(g) ⇾ 2BaO2(s)
2. Group 2 hydroxides
The Group 2 hydroxides are also bases and will neutralise an acid such as hydrochloric acid to form the metal salt and water.
Mg(OH)2(s) + 2HCl(aq) ⇾ MgCl2(aq) + H2O(l)
The solubility of the metal hydroxides in water increases as you descend the group. Magnesium hydroxide is sparingly soluble and a suspension can be used as an antacid (milk of magnesia). Calcium hydroxide is better known as lime water, reacting with carbon dioxide to give a cloudy calcium carbonate precipitate.
Ca(OH)2(aq) + CO2(g) ⇾ CaCO3(s) + H2O(l)
Barium hydroxide readily dissolves in water to form a strongly alkaline solution (the result of a high concentration of aqueous hydroxide ions).
Ba(OH)2(s) + aq ⇾ Ba2+(aq) + 2OH–(aq)
Although the Group 2 hydroxides can be prepared by reacting either the metal or the metal oxide with water (see above), neither approach works well for beryllium which is best prepared by reacting beryllium chloride with an alkali such as sodium hydroxide.
BeCl2(s) + 2NaOH(aq) ⇾ Be(OH)2(s) + 2NaCl(aq)
3. Thermal stability of Group 2 compounds
The thermal stability of Group 2 carbonates, hydroxides and nitrates increases as you descend the group – we have to heat barium compounds to a higher temperature before they decompose compared with magnesium compounds.
SrCO3(s) ⇾ SrO(s) + CO2(g)
2Ca(NO3)2(s) ⇾ 2CaO(s) + 4NO(g) + 3O2(g)
Mg(OH)2(s) ⇾ MgO(s) + H2O(l)
| Decomposition temperature of MCO3 / °C | Decomposition temperature of M(NO3)2 / °C | |
|---|---|---|
| Mg | 400 | 450 |
| Ca | 900 | 575 |
| Sr | 1280 | 635 |
| Ba | 1360 | 675 |
The temperature at which a Group 2 compound decomposes spontaneously is a reflection of how endothermic the enthalpy change of reaction is. The enthalpy change for a decomposition reaction, ΔrH, mostly depends on the difference between the strength of the ionic bonds in the reactant lattice (the metal carbonate / nitrate / hydroxide) and the strength of these bonds in the product lattice (the metal oxide). This ‘strength’ can be calculated theoretically and it is called a lattice enthalpy.
MgO(s) ⇾ Mg2+(g) + O2-(g) ΔLEH⦵ = +3889 kJ mol-1
BaO(s) ⇾ Ba2+(g) + O2-(g) ΔLEH⦵ = +3152 kJ mol-1
Lattice enthalpies for individual Group 2 compounds become less endothermic as we descend the group as we can see above, and they are predominantly governed by the size of the ions in the lattice – the smaller the ions, the closer they will be in the lattice and the stronger the electrostatic attractions between them. This equates to a more positive / more endothermic lattice enthalpy as more energy is needed to break up the lattice.
We find that the lattice enthalpies for Group 2 carbonates, for example, are less endothermic than those for the Group 2 oxides, for exactly the same reasons – oxide ions are small and have a high charge density compared with carbonate ions.
MgO(s) ⇾ Mg2+(g) + O2-(g) ΔLEH⦵ = +3889 kJ mol-1
MgCO3(s) ⇾ Mg2+(g) + CO32-(g) ΔLEH⦵ = +3123 kJ mol-1
However, it is the trend in the overall enthalpy of reaction for thermal decomposition that we are interested in, and as we descend the group this becomes more endothermic (more positive).
Partly this is due to the difference in the lattices enthalpies of the metal carbonates / nitrates / hydroxides and the metal oxides, and partly due to the enthalpy change associated with the decomposition of the carbonate / nitrate / hydroxide anion itself (see extension below).
The trend is that magnesium carbonate, magnesium nitrate and magnesium hydroxide are the least thermally stable in Group 2, and their barium counterparts are the most thermally stable.
Extension …
For a complete picture we also need to consider the enthalpy change for the anion (carbonate / nitrate / hydroxide) decomposing into an oxide ion and gaseous products.
We could assume that it is independent of whether the metal cation is magnesium, calcium, strontium or barium, but there is an argument that in magnesium compounds, the high charge density of the Mg2+ ion polarises larger anions such as carbonate and nitrate in a similar way to Be2+. The strain this places on the anion means that it takes a little less energy for it to disproportionate into an oxide ion and gas (either carbon dioxide, nitrogen oxide or oxygen).
4. Solubility of Group 2 sulphates
The solubility of the Group 2 sulphates decreases as you descend the group – magnesium sulphate is very soluble, barium sulphate is very definitely insoluble. This is related to the size of the cation since small cations such as Mg2+ with a higher charge density are more easily hydrated than large cations (Ba2+). You can find out everything you need to know about solubility on this page 😎.
Exam question practice
- Write balanced equations for the following reactions:
(a) burning calcium in excess air
(b) the reaction between strontium hydroxide and dilute nitric acid, HNO3
(c) the reaction between barium oxide and water
- When magnesium burns in air the main product of the reaction is magnesium oxide but magnesium nitride, Mg3N2, is also formed.
(a) Explain why these two products are formed in the reaction.
(b) Given that magnesium nitride dissolves in water to form magnesium hydroxide, suggest how a pure sample of magnesium oxide could be obtained from the mixture of products.
- Hydrated crystals of magnesium chloride, MgCl2•𝑥H2O can be prepared from the reaction between magnesium hydroxide and hydrochloric acid. Given that a sample of the chloride contained 16.5% magnesium by mass, determine how many water of crystallisation are present.
Answers
- (a) 2Ca(s) + O2(g) ⇾ 2CaO(s)
(b) Sr(OH)2(s) + 2HNO3(aq) ⇾ Sr(NO3)2(aq) + 2H2O(l)
(c) BaO(s) + H2O(l) ⇾ Ba(OH)2(aq)
- (a) Magnesium burns in the oxygen in the air but due to the very high temperature of this reaction the magnesium also reacts with nitrogen in the air.
(b) Magnesium oxide and magnesium nitride both dissolve in water to give magnesium hydroxide. Heating magnesium hydroxide strongly will cause it to thermally decompose giving a pure sample of magnesium oxide.
- If the sample is 16.5% Mg by mass then in 100g of the chloride there will be 16.5g Mg.
no. of mol of Mg = mass / Mr = 16.5 / 24 = 0.69mol
ratio of Mg:Cl in the compound is 1:2 so there must be 1.38 mol of Cl present
mass of Cl = no. of mol x Mr = 1.38 x 35.5 = 48.99g
mass of water = 100 – (18.5 + 48.99) = 32.51g
no. of mol of water = mass / Mr = 32.51 / 18 = 1.81mol so 𝓍 = 2 water of crystallisation