How Reactive Are the Group 2 Metals?

1Same family, very different speeds

Drop a clean coil of magnesium ribbon into a beaker of cold water and you may wait a whole lesson to collect a few bubbles. Drop a granule of calcium into the same water and it fizzes at once, the water turns milky, and the granule is gone in a minute or two. Barium reacts faster still. Yet hot magnesium burns fiercely in steam. In this lesson you will compare the reactions of the Group 2 metals with oxygen, water, steam, acids and chlorine, and explain the order of reactivity from the energy needed to form the ${\mathrm{M}}^{2+}\backslash mathrm\{M^\{2+\}\}$ ion.

2With oxygen and nitrogen

All the metals burn when heated in air or oxygen to form the white ionic oxide MO:$$2\mathrm{M}(\mathrm{s})+{\mathrm{O}}_2(\mathrm{g})\to 2\mathrm{M}\mathrm{O}(\mathrm{s})\backslash mathrm\{2M(s)\; +\; O\_2(g)\; \backslash rightarrow\; 2MO(s)\}$$Magnesium burns with a dazzling white light, calcium with a brick-red flame, strontium crimson and barium pale green. Beryllium burns only as a fine powder, because a thin oxide layer protects the solid. Barium heated in excess oxygen also forms barium peroxide, $\mathrm{B}\mathrm{a}{\mathrm{O}}_2\backslash mathrm\{BaO\_2\}$, which contains the peroxide ion ${\mathrm{O}}_2^{2-}\backslash mathrm\{O\_2^\{2-\}\}$; the large $\mathrm{B}{\mathrm{a}}^{2+}\backslash mathrm\{Ba^\{2+\}\}$ ion can stabilise this large anion.

When magnesium burns in air, some of it also combines with nitrogen to form magnesium nitride, a pale yellow solid:$$3\mathrm{M}\mathrm{g}(\mathrm{s})+{\mathrm{N}}_2(\mathrm{g})\to \mathrm{M}{\mathrm{g}}_3{\mathrm{N}}_2(\mathrm{s})\backslash mathrm\{3Mg(s)\; +\; N\_2(g)\; \backslash rightarrow\; Mg\_3N\_2(s)\}$$Water turns the nitride into magnesium hydroxide and ammonia, which you can smell or detect with damp red litmus:$$\mathrm{M}{\mathrm{g}}_3{\mathrm{N}}_2(\mathrm{s})+6{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to 3\mathrm{M}\mathrm{g}(\mathrm{O}\mathrm{H}{)}_2(\mathrm{s})+2\mathrm{N}{\mathrm{H}}_3(\mathrm{g})\backslash mathrm\{Mg\_3N\_2(s)\; +\; 6H\_2O(l)\; \backslash rightarrow\; 3Mg(OH)\_2(s)\; +\; 2NH\_3(g)\}$$

3With water and with steam

Calcium, strontium and barium react with cold water, faster and faster down the group, giving the hydroxide and hydrogen:$$\mathrm{C}\mathrm{a}(\mathrm{s})+2{\mathrm{H}}_2\mathrm{O}(\mathrm{l})\to \mathrm{C}\mathrm{a}(\mathrm{O}\mathrm{H}{)}_2(\mathrm{a}\mathrm{q})+{\mathrm{H}}_2(\mathrm{g})\backslash mathrm\{Ca(s)\; +\; 2H\_2O(l)\; \backslash rightarrow\; Ca(OH)\_2(aq)\; +\; H\_2(g)\}$$Calcium hydroxide is only slightly soluble, so some of it appears as a white suspension that makes the water milky. Barium hydroxide is much more soluble and the solution stays nearly clear (Lesson 3).

Magnesium reacts only very slowly with cold water, because the slightly soluble magnesium hydroxide coats the metal. Heated magnesium reacts vigorously with steam, glowing brightly and leaving white magnesium oxide, because any hydroxide would decompose at that temperature:$$\mathrm{M}\mathrm{g}(\mathrm{s})+{\mathrm{H}}_2\mathrm{O}(\mathrm{g})\to \mathrm{M}\mathrm{g}\mathrm{O}(\mathrm{s})+{\mathrm{H}}_2(\mathrm{g})\backslash mathrm\{Mg(s)\; +\; H\_2O(g)\; \backslash rightarrow\; MgO(s)\; +\; H\_2(g)\}$$Beryllium does not react with water or steam, even when red hot.

Figure 7.3: Reacting magnesium with steam (a teacher demonstration). Drawn by Ganzaa.

Safety for this demonstration. Use a hard-glass tube and safety screens, and wear eye protection: the tube can crack. Light the hydrogen only after the air has been driven out, and heat the magnesium, not the wet wool, so the tube does not fill with steam faster than the metal can use it. Do not look straight at burning magnesium; its light is intense.

4With dilute acids and with chlorine

All the metals react with dilute hydrochloric acid, giving a salt and hydrogen. Even beryllium reacts, slowly:$$\mathrm{M}\mathrm{g}(\mathrm{s})+2\mathrm{H}\mathrm{C}\mathrm{l}(\mathrm{a}\mathrm{q})\to \mathrm{M}\mathrm{g}\mathrm{C}{\mathrm{l}}_2(\mathrm{a}\mathrm{q})+{\mathrm{H}}_2(\mathrm{g})\backslash mathrm\{Mg(s)\; +\; 2HCl(aq)\; \backslash rightarrow\; MgCl\_2(aq)\; +\; H\_2(g)\}$$With dilute sulphuric acid, magnesium reacts steadily, but calcium, strontium and barium soon stop: their sulphates are only slightly soluble or insoluble and form a coat over the metal. Heated in chlorine, every metal forms its chloride, $\mathrm{M}(\mathrm{s})+\mathrm{C}{\mathrm{l}}_2(\mathrm{g})\to \mathrm{M}\mathrm{C}{\mathrm{l}}_2(\mathrm{s})\backslash mathrm\{M(s)\; +\; Cl\_2(g)\; \backslash rightarrow\; MCl\_2(s)\}$; $\mathrm{B}\mathrm{e}\mathrm{C}{\mathrm{l}}_2\backslash mathrm\{BeCl\_2\}$ is covalent and the others ionic.

Table 7.3: Reactions of the Group 2 metals

The last column gives the standard electrode potential, which you will meet properly in S5: the more negative it is, the more easily the metal gives up electrons in water. It becomes more negative down the group.

5Why reactivity increases down the group

To react, each atom must lose two electrons. The energy this costs, $I{E}_1+I{E}_{2}IE\_1\; +\; IE\_2$, falls by over 1100 kJ mol⁻¹ from beryllium to barium (Figure 7.4), because the outer electrons are further from the nucleus and better shielded. Less energy is needed to form the ion, so the reactions start more easily and go faster. (Energy is released when the ions form a lattice or are hydrated; this energy also gets smaller down the group, but by less than the ionisation energy falls.)

Figure 7.4: Energy needed to form ${\mathrm{M}}^{2+}(\mathrm{g})\backslash mathrm\{M^\{2+\}(g)\}$ from $\mathrm{M}(\mathrm{g})\backslash mathrm\{M(g)\}$, the sum of the first two ionisation energies (rounded). Data: CRC Handbook of Chemistry and Physics, via Wikipedia's ionisation energy data page.