Chemical Properties

Because of their large atomic size and very low ionization enthalpy, the alkali metals are the most reactive metals known, and reactivity itself increases further down the group.

1. Reactivity towards air Exposed to dry air, the metals tarnish as a surface layer of oxide forms; this oxide then reacts further with atmospheric moisture to give the hydroxide. Burnt in oxygen, each metal gives a characteristically different oxide: lithium gives only the normal oxide, sodium gives the peroxide, and the heavier members give superoxides (the bulky superoxide ion $O_2^-$ is stabilised only when paired with a correspondingly large cation such as $K^+$, $Rb^+$ or $Cs^+$).

   $4Li + O_2 \rightarrow 2Li_2O \quad (\text{oxide})$

   $2Na + O_2 \rightarrow Na_2O_2 \quad (\text{peroxide})$

   $M + O_2 \rightarrow MO_2 \quad (\text{superoxide}),\ M = K, Rb, Cs$

   In every one of these oxides the metal itself is in the +1 state; only the anion (oxide, peroxide or superoxide) changes down the group. Lithium alone also combines directly with atmospheric nitrogen to give the nitride, $Li_3N$. Because all the alkali metals react so readily with both air and water, they are stored under kerosene oil in the laboratory.
2. Reactivity towards water All alkali metals react with water to liberate hydrogen and form the metal hydroxide:

   $2M + 2H_2O \rightarrow 2M^+ + 2OH^- + H_2$

   Interestingly, lithium — despite having the most negative standard electrode potential in the group — reacts with water noticeably less violently than sodium, whose electrode potential is the least negative of the five. The explanation lies outside thermodynamics: lithium's very small ionic size gives it an enormous hydration energy, which affects reaction kinetics even though E° favours lithium overall. The heavier members (K, Rb, Cs) react explosively with water. All the alkali metals also react with other proton donors such as alcohols, gaseous ammonia and terminal alkynes.
3. Reactivity towards dihydrogen On heating to about 673 K (lithium needs a higher 1073 K), the alkali metals combine directly with hydrogen to give ionic hydrides, all of which are high-melting solids:

   $2M + H_2 \rightarrow 2M^+H^-$

4. Reactivity towards halogens The alkali metals react vigorously with the halogens to give ionic halides, $M^+X^-$. Lithium halides, however, are noticeably more covalent than the rest. This traces back to the very small $Li^+$ ion, which strongly polarises (distorts the electron cloud of) the neighbouring halide ion — an effect called polarisation. Since large, easily-distorted anions polarise most readily, lithium iodide (with the largest halide ion) is the most covalent of the lithium halides.
5. Reducing nature All the alkali metals are powerful reducing agents; among them lithium is the strongest and sodium the weakest, judged by standard reduction potential (Table 10.1). That overall E° is really the net result of three separate steps:

   $M(s) \rightarrow M(g) \quad \text{(sublimation enthalpy)}$

   $M(g) \rightarrow M^+(g) + e^- \quad \text{(ionization enthalpy)}$

   $M^+(g) + H_2O \rightarrow M^+(aq) \quad \text{(hydration enthalpy)}$

   Lithium's exceptionally small ionic size gives it, by a wide margin, the largest hydration enthalpy of the three terms — and this large hydration enthalpy dominates the overall energetics, producing lithium's strongly negative E° and hence its position as the strongest reducing agent in the group.
6. Solutions in liquid ammonia …