Metal hydroxide tests and writing balanced equations

Chemical analysis • Identification of ions

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Write the ionic equation for aluminium hydroxide precipitate.

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Al3+ + 3OH- -> Al(OH)3(s).

Key concepts

What you'll likely be quizzed about

Principle of the hydroxide precipitation test

Addition of hydroxide ions (commonly as NaOH) to a solution of metal ions causes formation of a metal hydroxide if that hydroxide is insoluble. Precipitation occurs when the ionic product exceeds the solubility product, so a visible solid forms from the ionic mixture. The ionic equation representation isolates the reacting ions and the precipitate species for clarity.

Common precipitate colours and identifications

Copper(II) ions produce a blue precipitate of copper(II) hydroxide. Iron(II) ions produce a pale green precipitate of iron(II) hydroxide. Iron(III) ions produce a brown or reddish-brown precipitate of iron(III) hydroxide. Aluminium, calcium and magnesium ions produce white precipitates that require further testing to distinguish.

Amphoteric hydroxides and behaviour with excess NaOH

Certain metal hydroxides are amphoteric and dissolve in excess hydroxide to form soluble complex ions. Aluminium hydroxide and zinc hydroxide both dissolve in excess NaOH to give aluminate and zincate ions respectively. Copper(II), calcium and magnesium hydroxides do not dissolve in excess hydroxide under normal test conditions.

Test for ammonium ions

Addition of NaOH to an ammonium-containing solution followed by gentle warming produces ammonia gas rather than a solid hydroxide. The gas turns damp red litmus paper blue and gives a characteristic pungent smell. The ionic equation for this reaction is NH4+ + OH- -> NH3(g) + H2O.

Writing balanced equations for precipitate formation

Balanced ionic equations for hydroxide precipitates combine the metal cation and hydroxide anions in stoichiometric ratios to form the solid hydroxide. Example form: M^n+ + nOH- -> M(OH)n(s). Full equations with common salts show spectator ions and ensure mass and charge balance.

Limiting factors and interferences

Oxidation, concentration, and other ions affect test outcomes. Iron(II) precipitates oxidise to iron(III) species on exposure to air, changing the colour. Carbonates in the sample produce additional precipitates or gas evolution. Very low ion concentrations give weak or no visible precipitate, and coloured solutions can mask precipitate colour.

Key notes

Important points to keep in mind

Add sodium hydroxide dropwise and observe precipitate colour immediately.

Test with excess NaOH to detect amphoteric hydroxides (Al3+, Zn2+ dissolve).

Warm the mixture to test for ammonium ions; look for ammonia gas and litmus change.

Write ionic equations to show only species that change during precipitation.

Use full equations with spectator ions to demonstrate mass and charge balance.

Consider oxidation: Fe2+ precipitates darken on exposure to air, indicating Fe3+ formation.

Low concentrations or colored solutions can mask weak white precipitates.

Carbonate contamination produces additional precipitates and interferes with interpretation.

Differentiate white precipitates (Al3+, Ca2+, Mg2+) by testing solubility in excess NaOH and using further tests.

State symbols clarify equations: (aq) for dissolved ions, (s) for precipitates, (g) for ammonia.