Grade 12

Alkali metals, alkaline earth metals, aluminium

Understand alkali and alkaline-earth trends, flame colours, hardness treatment and aluminium extraction through ionic chemistry, electrochemistry and bonding.

1 Lower secondary 2 Upper secondary 3 Undergraduate 4 Advanced 5 Research

IntuitionMetals that readily give up electrons

Alkali metals have one outer s electron; alkaline-earth metals have two. Both form largely ionic compounds, but their physical properties and reactions vary systematically with charge, size and oxide chemistry. Aluminium is different: its strong oxide film makes a reactive metal appear durable.

Rotate sodium’s bcc lattice: each atom has eight nearest neighbours. Adjust the number of cells, atom size, unit-cell edges and vacancies.

The selected metal is sodium (metal 1); amount 0.8 controls the displayed sample amount. Flame colour is characteristic, not a quantitative concentration reading.

SchoolAlkali metals and their compounds

Alkali metals are soft, low-density metals with relatively low melting points and bcc structures at ordinary conditions. Reactivity with water increases from Li to Cs: hydrogen and hydroxide form, with heat released. Store reactive specimens under suitable oil, isolated from water and ignition sources.

\ce{2M + 2H2O -> 2MOH + H2}

Lithium burns mainly to Li₂O; sodium commonly forms Na₂O₂ in oxygen, while potassium can form KO₂ under suitable oxygen-rich conditions. Industrial chlor-alkali electrolysis of brine produces chlorine, hydrogen and sodium hydroxide. The Solvay process makes sodium carbonate; sodium hydrogen carbonate is a useful intermediate and leavening agent.

SchoolAlkaline-earth metals

Group 2 metals form M²⁺ ions. Magnesium burns with a brilliant white light; calcium reacts with cold water more slowly than the heavier alkaline-earth metals. Limestone, quicklime and slaked lime form a useful cycle. Temporary hardness from hydrogencarbonates can be reduced by boiling; carbonate precipitation can remove Ca²⁺/Mg²⁺ permanently.

\ce{CaCO3 ->[\Delta] CaO + CO2}

\ce{CaO + H2O -> Ca(OH)2}

Example: Lime-soda softening

A sample contains 1.00 mmol L⁻¹ Ca(HCO₃)₂ and 1.00 mmol L⁻¹ CaCl₂. What ideal reagent amounts remove these hardness ions from 1.00 L?

Solution

Temporary hardness: $\ce{Ca(HCO3)2 + Ca(OH)2 -> 2CaCO3(s) + 2H2O}$ , so 1.00 mmol Ca(OH)₂ is required. Permanent hardness: $\ce{CaCl2 + Na2CO3 -> CaCO3(s) + 2NaCl}$ , so 1.00 mmol Na₂CO₃ is required. Real treatment must account for other ions, pH and practical excess.

Flame colours are useful qualitative clues: Li red, Na yellow, K lilac, Sr red and Ba green. The colour arises from element-specific electronic transitions; mixtures and contamination can obscure it.

Selected physical data near 1 bar
Metal	Melting point / °C	Density / g cm⁻³	First ionization energy / kJ mol⁻¹
Li	180.5	0.534	520.2
Na	97.8	0.968	495.8
K	63.4	0.862	418.8
Mg	650	1.738	737.7
Ca	842	1.55	589.8

Characteristic flame colours
Element	Observed colour
Li	Crimson red
Na	Yellow
K	Lilac
Sr	Crimson red
Ba	Green

SchoolAluminium: passivation and electrolysis

A thin, adherent Al₂O₃ film passivates aluminium in air. Aluminium hydroxide is amphoteric: it dissolves in acid and in excess strong base. The Hall–Héroult process electrolyses alumina dissolved in molten cryolite (Na₃AlF₆) at about 950 °C; carbon anodes are consumed as oxygen-containing products form. Thermite reduction of iron(III) oxide by aluminium is highly exothermic.

\ce{Al(OH)3 + 3H+ -> Al3+ + 3H2O}

\ce{Al(OH)3 + OH- -> [Al(OH)4]-}

Example: Faraday-law aluminium yield

An ideal cell passes 10.0 A for 30.0 min. Find the Al mass at 100% current efficiency.

Solution

For $\ce{Al^3+ + 3e- -> Al}$ , $n=3$ and $Q=It=1.80\times10^4\,C$ . Thus $m=MIt/(nF)=(26.98)(10.0)(1800)/(3\times96485)=1.68\,g$ Al. Actual efficiency is below 100%.

\ce{Fe2O3 + 2Al -> 2Fe + Al2O3}

UndergraduateWhy properties change down the groups

First ionization energy generally falls down a group because the valence electron occupies a shell farther from the nucleus and is increasingly shielded, outweighing the increased nuclear charge. Lithium and magnesium show a diagonal relationship: similar ionic size/charge density gives some comparable chemistry (for example, nitrides and less soluble carbonates), though they are not identical. Beryllium and aluminium also share covalent, amphoteric tendencies.

The thermal stability of group-2 carbonates increases down the group: larger M²⁺ ions have lower charge density and polarise CO₃²⁻ less, so decomposition is less favored. Solubility trends are not read from lattice energy alone: dissolution balances lattice separation against hydration. Down group 2, hydroxide solubility generally increases, while sulfate solubility generally decreases as hydration becomes less favorable relative to lattice stabilization.

References

Inorganic Chemistry, 5th edition · Catherine E. Housecroft and Alan G. Sharpe, 2018
Chemistry of the Elements, 2nd edition · N. N. Greenwood and A. Earnshaw, 1997