The s-Block Elements: Alkali and Alkaline Earth Metals

Groups 1 and 2 together make up the s-block, sharing the defining feature of having their outermost electron(s) in an s orbital only, ns1 for the alkali metals and ns2 for the alkaline earth metals, and both groups build their entire chemistry around losing exactly that many electrons to reach a stable, noble-gas-like configuration. This single valence-electron count difference (one versus two) explains the two groups' most basic contrasts directly: alkali metals show only the +1 oxidation state and form largely ionic, water-soluble compounds; alkaline earth metals show the +2 state, and their generally smaller size and higher charge give their compounds noticeably higher melting points, higher lattice enthalpies, and, for several key compounds, markedly lower solubility than their group 1 counterparts. Both groups are strong reducing agents, readily giving up their valence electrons to whatever oxidising agent is available, and both groups' reactivity increases going down the group, as increasing atomic size makes that outermost electron progressively easier to remove.

Alkali metals are genuinely unusual among metals: they are soft enough to be cut with an ordinary knife, and lithium, sodium, and potassium are all actually less dense than water, floating on its surface even as they react with it. All alkali metals have low melting points (falling further down the group, as the metallic bond weakens with increasing atomic size and correspondingly poorer orbital overlap) and characteristic flame colours used in flame tests, lithium crimson-red, sodium persistent yellow, and potassium lilac, each colour arising from electrons excited by flame heat falling back down through the atom's own specific energy-level spacing and releasing that exact wavelength of visible light. Because alkali metals are so reactive toward both oxygen and moisture in ordinary air, they must be stored under an inert liquid such as kerosene or mineral oil to prevent uncontrolled reaction.

Reactivity of the alkali metals with water increases sharply down the group: lithium reacts steadily but without catching fire, sodium reacts vigorously enough to melt into a fizzing ball and often ignite, and potassium reacts violently enough to ignite immediately and essentially always catch fire. This trend appears, at first glance, to contradict standard electrode potential data, since lithium actually has the most negative (most favourable, thermodynamically) standard reduction potential of the whole group, which would suggest it should react most vigorously, not least. The resolution is that standard electrode potentials describe overall thermodynamic favourability, not the actual rate of reaction, and lithium's very small size gives it an unusually high enthalpy of hydration that, combined with a very high lattice/atomisation enthalpy for the solid metal itself, actually makes its reaction proceed more slowly in practice than potassium's much larger, more loosely-held atoms allow, a genuine, frequently tested example of the gap between thermodynamic favourability and kinetic reaction rate.

Lithium is genuinely the odd one out within its own group, behaving noticeably differently from sodium, potassium, and the rest, for reasons that trace back to its unusually small size and correspondingly high charge density (polarising power): lithium is the only alkali metal to react directly with atmospheric nitrogen to form a nitride (6Li + N2 to 2Li3N), its compounds show noticeably more covalent character than the other alkali metals' largely ionic compounds, and lithium chloride, unlike the other alkali halides, is soluble in organic solvents like ethanol, another direct sign of that same covalent character. These anomalies are specific examples of a broader pattern called the diagonal relationship, in which the first member of a group resembles the element diagonally below-and-across from it in the next group more closely than it resembles its own heavier group-mates, and here specifically, lithium's small size and high charge density give it chemistry that resembles magnesium (group 2) more closely than it resembles sodium or potassium.

Sodium hydroxide, NaOH, is manufactured industrially by the chlor-alkali (Castner-Kellner) process, the electrolysis of concentrated brine (aqueous sodium chloride), producing NaOH, chlorine gas, and hydrogen gas simultaneously as three separately valuable products. Sodium carbonate, Na2CO3 (washing soda), is manufactured by the Solvay process, an ingenious sequence starting from cheap brine and limestone that ultimately produces sodium carbonate while regenerating and recycling most of its own ammonia catalyst, making it remarkably efficient at industrial scale. Sodium hydrogencarbonate, NaHCO3 (baking soda), decomposes on heating to release carbon dioxide gas, 2NaHCO3 to Na2CO3 + H2O + CO2, exactly the reaction that makes baked goods rise, and it is also mildly basic in solution, giving it its familiar use as a stomach-acid antacid.

Group 2, the alkaline earth metals, show a family resemblance to group 1 but with several systematic differences directly traceable to their ns2 configuration and correspondingly higher nuclear charge for a similar size: they are generally harder, denser, and have higher melting points than the corresponding alkali metals in the same period, since two valence electrons per atom (rather than one) are available for metallic bonding, producing a stronger metallic bond overall. Their first and second ionisation enthalpies combined are correspondingly higher than an alkali metal's single ionisation enthalpy, yet group 2 metals remain excellent reducing agents overall, readily forming +2 ions, and reactivity still increases down the group for the same underlying reason as group 1, growing atomic size making the outer electrons progressively easier to remove.

Beryllium is the group 2 counterpart of lithium's anomaly, an even more extreme case, since beryllium is smaller and has an even higher charge density than any other group 2 element: beryllium compounds are predominantly covalent rather than ionic (BeCl2, for instance, exists as a covalent, chain-polymeric solid rather than a simple ionic lattice), beryllium's own hydroxide, Be(OH)2, is genuinely amphoteric (dissolving in both acids and strong alkalis, unlike any other group 2 hydroxide), and beryllium never shows the coordination number of six that the rest of group 2 commonly adopts, being restricted instead to a maximum coordination number of four. These anomalies mirror lithium's own diagonal relationship almost exactly: beryllium's small size and high charge density give it a genuine diagonal relationship with aluminium (group 13), resembling aluminium's amphoteric, largely covalent chemistry more closely than it resembles magnesium, calcium, or the rest of its own group.

Group 2 hydroxide solubility in water increases going down the group, from essentially insoluble Be(OH)2 and Mg(OH)2 through to genuinely soluble Ca(OH)2, Sr(OH)2, and Ba(OH)2, while group 2 sulphate solubility does exactly the opposite, decreasing down the group, from very soluble BeSO4 and MgSO4 through to essentially insoluble BaSO4, and both trends trace back to the very same underlying competition between lattice enthalpy and hydration enthalpy, just landing on opposite sides of that competition for two different anions. Both lattice enthalpy and hydration enthalpy fall as the cation grows larger down the group, but they fall at genuinely different rates depending on the anion's own size: for the small hydroxide ion, lattice enthalpy falls faster than hydration enthalpy as the cation grows, tipping the balance toward dissolving; for the considerably larger sulphate ion, hydration enthalpy instead falls faster than lattice enthalpy as the cation grows, tipping the balance the opposite way, toward staying locked in the solid lattice instead.

Group 1 and group 2 carbonates and nitrates both become thermally more stable moving down their respective groups, and the explanation rests on a single idea called the polarising power of the cation: a smaller, more highly charged cation distorts (polarises) the large, diffuse carbonate or nitrate anion's electron cloud more strongly, weakening the anion's own internal C-O or N-O bonding and making it correspondingly easier for the anion to break down on heating into a smaller oxide ion plus CO2 or into oxide plus nitrogen oxides. Since cations grow larger and less polarising moving down either group, decomposition temperature climbs steadily down the group in both families; magnesium carbonate decomposes at a noticeably lower temperature than barium carbonate, for exactly this reason, and lithium is again the specific anomaly worth remembering within group 1, being the only alkali metal carbonate that decomposes on heating at all, since lithium's own unusually small, highly polarising cation behaves more like a group 2 cation in this specific respect too, one more direct expression of its diagonal relationship with magnesium.

Calcium carbonate, CaCO3 (limestone), decomposes on strong heating in a lime kiln to give calcium oxide (quicklime) and carbon dioxide, CaCO3 to CaO + CO2, and quicklime reacts vigorously and exothermically with water to give calcium hydroxide (slaked lime), CaO + H2O to Ca(OH)2, a cheap, strong base used throughout construction and agriculture. Plaster of Paris, calcium sulphate hemihydrate (CaSO4.½H2O), is made by carefully heating gypsum (CaSO4.2H2O) to drive off three-quarters of its water of crystallisation, and it sets hard again when mixed with water because that same water is simply re-absorbed, reforming the original interlocking gypsum crystal structure, which is exactly why Plaster of Paris is so useful for casts and moulds. Both groups also play genuinely essential biological roles: sodium and potassium ions maintain the electrochemical gradient across nerve cell membranes that makes nerve impulse transmission and muscle contraction possible, magnesium sits at the very centre of the chlorophyll molecule that drives photosynthesis, and calcium is the primary structural component of bones and teeth while also triggering blood clotting and muscle contraction directly.