Hydrogen and the s-block

Drop a small piece of lithium into water and it floats, fizzing quietly; sodium melts into a silver bead that skates around; potassium bursts into a lilac flame at once; caesium simply explodes. Almost every textbook says that “the alkali metals get more reactive down the group”. Yet if you calculate how much energy each reaction releases, the answer is a surprise.

Hydrogen has a single electron, so it can lose it like an alkali metal to become $\ce{H+}$, gain one like a halogen to become $\ce{H-}$, or share it. It has three isotopes: ordinary protium $\ce{^1H}$, deuterium $\ce{^2H}$ at about 0.0156 % of the hydrogen in seawater (roughly 1 in 6400; D — heavy water $\ce{D2O}$ moderates some nuclear reactors), and radioactive tritium $\ce{^3H}$ (T, half-life 12.32 years).

Hydrides fall into three families according to the partner element:

• Ionic hydrides form with the most electropositive s-block metals and contain $\ce{H-}$: NaH, $\ce{CaH2}$. Hydride is a very strong base and reducing agent — $\ce{CaH2} + \ce{2H2O} \mathrel{\ce{->}} \ce{Ca(OH)2} + \ce{2H2}$ — so $\ce{CaH2}$ serves as a drying agent and a portable source of hydrogen.
• Covalent hydrides form with the p-block non-metals: $\ce{CH4}$, $\ce{NH3}$, $\ce{H2O}$, HF — their properties are set by polarity and hydrogen bonding.
• Metallic (interstitial) hydrides: many d-block metals soak hydrogen atoms into the holes of their lattice. Palladium can take up hundreds of times its own volume of hydrogen at room temperature, with a composition that varies continuously ($\ce{PdH_x}$).

A group-1 metal reacts with water as $\ce{M(s)} + \ce{H2O(l)} \mathrel{\ce{->}} \ce{M+(aq)} + \ce{OH-(aq)} + \ce{1/2H2(g)}$. Hess’s law lets us split the metal’s side into three steps: pull the metal apart into gaseous atoms (atomisation, endothermic), strip off the valence electron (ionisation, endothermic), and let water molecules crowd round the gaseous ion (hydration, exothermic). The water side ($\ce{H2O} + \ce{e-} \mathrel{\ce{->}} \ce{OH-} + \ce{1/2H2}$) is the same for every metal.

From lithium to caesium the “cost” of atomising and ionising falls from 680 to 452 kJ/mol — the metals give up their electron ever more easily — but the ions also grow, so the hydration “payback” shrinks from −519 to −271 kJ/mol. The two nearly cancel, so ΔH° and ΔG° hardly change down the group, and lithium even comes out slightly ahead: the tiny $\ce{Li+}$ is gripped hardest by water.

So what sets the violence? The rate. The metal has to hand electrons to water at its surface, and how easily it does so follows the cost of giving up an electron — which falls steadily from lithium to caesium. Moreover, everything below lithium melts below 100 °C (sodium 98 °C, potassium 63 °C, caesium 28 °C), and the heat of reaction (about 184 kJ per mole for sodium, some forty times what it takes to warm and melt it) melts the metal. A molten bead keeps renewing and enlarging its surface, the reaction accelerates, and the hydrogen ignites. Lithium melts only at 181 °C and reacts slowly, so water carries its heat away in time.

Down group 1 the metallic radius grows from 157 to 272 pm, the first ionisation energy falls, metallic bonding weakens and the melting point drops from 181 °C to 28 °C; density generally rises, except that potassium is lighter than sodium (its atomic volume grows faster than its mass). In group 2 two valence electrons join the metallic bonding, so melting points and hardness are far higher.

The salts of these elements colour flames: lithium crimson, sodium yellow, potassium lilac, calcium brick-red, strontium crimson, barium green. Sodium and lithium colours come from atomic lines (sodium’s 589 nm doublet); calcium, strontium and barium colours come mainly from molecules formed in the flame such as CaOH, SrOH and BaCl — see them resolved in the flame and discharge lab.

The solubilities of group-2 compounds run in opposite directions: hydroxides become more soluble down the group, sulfates less soluble. For the small $\ce{OH-}$ ion the lattice enthalpy falls quickly as the cation grows, outpacing the fall in hydration enthalpy; for the large $\ce{SO4^2-}$ ion the lattice enthalpy is dominated by the big anion and changes little, while the cation’s hydration enthalpy keeps falling. Everyday chemistry uses both trends: $\ce{Mg(OH)2}$ is the antacid “milk of magnesia”, limewater is saturated $\ce{Ca(OH)2}$, and a “barium meal” is almost insoluble $\ce{BaSO4}$.

In many ways lithium resembles magnesium, diagonally below and to the right, more than its own group-mate sodium; beryllium likewise resembles aluminium. The reason is a similar combination of size and charge: $\ce{Li+}$ (76 pm) and $\ce{Mg^2+}$ (72 pm) are almost the same size, while $\ce{Be^2+}$ (45 pm) and $\ce{Al^3+}$ (54 pm) have similar charge-to-radius ratios and polarise their neighbours strongly.

• Lithium and magnesium both burn in nitrogen to nitrides ($\ce{Li3N}$, $\ce{Mg3N2}$); their carbonates decompose on heating ($\ce{Li2CO3} \mathrel{\ce{->}} \ce{Li2O} + \ce{CO2}$, while the other alkali carbonates are stable); their fluorides, carbonates and phosphates are sparingly soluble.
• Beryllium and aluminium both have amphoteric hydroxides; their chlorides are covalent Lewis acids that satisfy their electron-poor centres by bridging (chains of $\ce{(BeCl2)_n}$, dimeric $\ce{Al2Cl6}$); and both are protected by dense oxide films that even concentrated nitric acid cannot breach.