s-Block Elements
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Direct answer
Group 1 metals are soft enough to cut with a knife, float on water (density under 1 g per cm³ for Li, Na, K) and paint flames in element-specific colours — lithium crimson, sodium golden yellow, potassium lilac, rubidium red-violet, caesium blue. Burning in excess air sorts them by size too: lithium stops at Li2O, sodium at the peroxide Na2O2, and the heavier three at superoxides MO2, whose large anions sit comfortably in the large cations' lattices. Lithium walks its own path, behaving like magnesium through the diagonal relationship, while the group 2 story is written by falling sulphate solubility — BaSO4 so insoluble that it is safely swallowed for X-rays.
What you must remember
- Flame colours: Li crimson red, Na golden yellow, K lilac (view through blue cobalt glass), Rb red-violet, Cs blue; beryllium and magnesium give no colour.
- Sodium and potassium are stored under kerosene; lithium floats even on kerosene, so it is kept wrapped in paraffin wax.
- Superoxide KO2 is paramagnetic and used in breathing masks and submarines, absorbing CO2 while releasing O2.
- Alkali hydroxides and carbonates grow more soluble down the group; Li2CO3 is the odd one that decomposes on heating, like MgCO3.
- Group 2 trends run opposite for sulphates: solubility falls from MgSO4 to BaSO4, while hydroxide solubility rises — Ca(OH)2 to Ba(OH)2.
- BeO is amphoteric (reacts with acid and alkali); the remaining oxides of both groups are basic.
- Solvay process cannot make potassium carbonate because KHCO3 is too soluble (and KCl too pricey) for the ammonia-recovery loop to close.
- Sodium carbonate is washing soda (Na2CO3·10H2O), used in glass making; NaOH is caustic soda, made by electrolysis of brine.
Reasoning with size: oxides and solubility
Why does potassium burn to the superoxide while lithium stops at the ordinary oxide? Lattice energy logic: a big cation cannot press tightly onto a small oxide ion, so the lattice of K2O is poorly stabilised; the peroxide and superoxide ions are progressively larger, and a large anion in a large cation's lattice maximises lattice energy. Sodium sits in the middle and settles on the peroxide. The same size-matching principle, run against hydration energy, decodes group 2 sulphates: down the group the cation grows, its hydration enthalpy falls steeply (charge unchanged, radius increasing), while the lattice energy with the big SO4^2− barely changes — the balance tips from dissolved (magnesium) to insoluble (barium). Reverse the anion to hydroxide, small enough that lattice energy falls faster than hydration enthalpy down the group, and the solubility trend flips upward. Two opposite trends from one competition — that is the chapter's intellectual core and the reason exam questions keep returning to it.
Flame tables and lithium's exceptions
The flame-colour table is cheap marks only if the pairs stay separated: rubidium red-violet versus caesium blue, and potassium observed through cobalt glass to filter sodium's overwhelming yellow contamination — practical details the exam likes to quote. Second, lithium's quirks: it forms the normal oxide and a stable nitride Li3N directly (like magnesium), its carbonate and nitrate decompose on heating unlike its group siblings', and LiF is sparingly soluble despite the small ion — each fact is a standalone question. Third, sodium hydroxide's manufacture: chlor-alkali electrolysis of brine gives NaOH, Cl2 and H2 together — students remember the NaOH and forget the two gases. Fourth, the biological pair: the sodium-potassium pump keeps K+ concentrated inside cells and Na+ outside, transmitting nerve impulses, while some 99 per cent of body calcium sits in bones and teeth and magnesium sits at the centre of the chlorophyll porphyrin. Finally, quicklime (CaO), slaked with water to Ca(OH)2, and bleaching powder from Ca(OH)2 and chlorine, are the group's industrial one-two that statement questions recycle.
Frequently asked questions
Why is lithium stored wrapped in paraffin wax?
Lithium is the lightest metal and floats on kerosene, exposing it to air anyway, so it is sealed in wax instead of the kerosene bath used for sodium and potassium.
Which alkali metal forms a superoxide, and what is special about it?
Potassium (and Rb, Cs) burn in excess air to KO2, a paramagnetic superoxide that absorbs CO2 and releases oxygen — used in breathing apparatus.
Why does the solubility of group 2 sulphates decrease down the group?
Hydration enthalpy falls faster than lattice energy as the cation grows, so the dissolution balance shifts against solution from MgSO4 to BaSO4.
Which s-block oxides are amphoteric?
Beryllium oxide (and to a lesser degree, beryllium hydroxide) — they dissolve in both acids and alkalis; the rest of both groups form basic oxides.
Why can't potassium carbonate be made by the Solvay process?
Potassium bicarbonate is far more soluble than sodium bicarbonate, so it never precipitates from the ammoniated brine, breaking the process's filtration step.