2_3 Reactions of Alkali and Alkaline Earth Metals with Water, Oxygen and Halogens

With one or two loosely held outer electrons, s-block metals are powerful reducing agents, meaning they give electrons to whatever they react with. Their chemistry is mostly a set of variations on one theme: the metal loses its s electrons and becomes

\mathrm{M^+}

or

\mathrm{M^{2+}}

. Group 1 has a very low first ionization energy and a huge second one, so it is limited to +1, while group 2 always forms +2 ions. What changes from element to element is how fast and how violently this happens.

Reactivity Trend Down the Groups

Reactivity increases down both groups. Each step down adds a shell, so the outer electron is further from the nucleus, better shielded and easier to remove. Potassium, rubidium and caesium react even with ice at very low temperatures, and caesium ignites on contact with air. Group 2 metals are less reactive than their group 1 neighbours because two electrons must be removed, and those electrons are held more tightly by a smaller atom with a larger nuclear charge.

Lithium has a twist. It is the strongest reducing agent of the group in aqueous solution, judged by electrode potential, because the tiny

\mathrm{Li^+}

ion releases a lot of energy when water molecules surround it. Yet it reacts with water more slowly than sodium. Electrode potential measures how far a reaction goes, not how fast. Lithium's higher melting point and higher ionization energy make the reaction slower to get going.

Reaction with Water

Alkali metals react with cold water to give the metal hydroxide and hydrogen gas. Sodium darts around on the surface, melting from the heat released, and potassium's reaction is hot enough to ignite the hydrogen with a lilac flame. The general equation is the same for all of them:

\mathrm{2M(s) + 2H_2O(l) \rightarrow 2MOH(aq) + H_2(g)} \qquad \text{(M = Li, Na, K, Rb, Cs)}

In group 2, beryllium does not react with water at all, because a protective oxide layer covers it. Magnesium barely reacts with cold water but burns in steam to form the oxide. Calcium, strontium and barium react steadily with cold water to form hydroxides, more vigorously further down the group.

\begin{aligned} \mathrm{Mg(s) + H_2O(g)} &\rightarrow \mathrm{MgO(s) + H_2(g)} \\ \mathrm{Ca(s) + 2H_2O(l)} &\rightarrow \mathrm{Ca(OH)_2(aq) + H_2(g)} \end{aligned}

Reaction with Oxygen: Oxides, Peroxides and Superoxides

Freshly cut alkali metals tarnish within seconds as an oxide forms. In moist air the oxide then picks up carbon dioxide to form a carbonate. When the metals burn in excess oxygen, the product depends on the size of the cation. Lithium forms the normal oxide (

\mathrm{O^{2-}}

), sodium forms the peroxide (

\mathrm{O_2^{2-}}

), and potassium, rubidium and caesium form superoxides (

\mathrm{O_2^-}

). Large oxygen anions are easily distorted, and only a large, weakly polarizing cation can sit next to them without breaking the O–O bond. Small

\mathrm{Li^+}

pulls them apart into simple oxide ions.

[Na].[Na].O=O>>[Na+].[Na+].[O-][O-]

\begin{aligned} \mathrm{4Li + O_2} &\rightarrow \mathrm{2Li_2O} \\ \mathrm{K + O_2} &\rightarrow \mathrm{KO_2} \end{aligned}

Group 2 metals burn to give normal oxides, MO. Magnesium burns with a dazzling white flame and also forms some magnesium nitride from the nitrogen in air. Barium, the largest, gives the peroxide

\mathrm{BaO_2}

when heated in oxygen at around 500 °C.

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Reaction with Halogens

All s-block metals combine directly with halogens to give ionic halides, MX in group 1 and

\mathrm{MX_2}

in group 2. Molten sodium burns in chlorine with a bright yellow flame. The heavier alkali metals react violently with every halogen.

[Na].[Na].ClCl>>[Na+].[Na+].[Cl-].[Cl-]

Hydrogen, Nitrogen and Sulphur

• Hydrogen: heated alkali metals form ionic hydrides such as NaH, which contain the hydride ion

  \mathrm{H^-}

  . Calcium, strontium and barium do the same, and magnesium needs high pressure and a catalyst. Hydrides react with water to release hydrogen,

  \mathrm{LiH + H_2O \rightarrow LiOH + H_2}

  , and serve as strong reducing agents.
• Nitrogen: lithium is the only alkali metal that forms a nitride, while all group 2 metals do so on heating. Nitrides react with water to give ammonia:

  \mathrm{Mg_3N_2 + 6H_2O \rightarrow 3Mg(OH)_2 + 2NH_3}

  .
• Sulphur: both groups combine with molten sulphur to give sulphides,

  \mathrm{M_2S}

  and MS.

Tips for remembering

• Oxide, peroxide, superoxide go down the group: Li₂O, Na₂O₂, KO₂. Bigger cations can stabilize bigger oxygen anions.
• Down the group, faster reaction: the outer electron gets further from the nucleus, so it is lost more easily and the metal reacts faster.
• Group 2 is one step tamer: two electrons have to go, so magnesium needs hot water or steam, while sodium fizzes on cold water.