7.6 - Tests for Cations

The flame test

• A cation is a positively charged ion.
• When a metal compound is heated in a flame, it gives off a characteristic colour that identifies the metal ion.

Carrying out the test

1. Clean the wire. Dip a nichrome or platinum wire loop in hydrochloric acid, then hold it in a blue Bunsen flame until it burns with no colour. Any residue left from an earlier test would mask the sample's colour. Hydrochloric acid is used because chloride ions give no flame colour.

2. Pick up the sample. Dip the clean wire loop into the solid sample so a little clings to it.

3. Hold the loop in the flame and record the colour. Use the roaring blue flame (air hole open), which is the hottest part and won't add any yellow colour of its own.

Flame colours

What the test cannot do

• A flame test only works reliably on a single metal ion.
• If a sample contains a mixture of ions, the colours blend and a strong colour masks a weaker one (e.g. sodium's yellow completely masks potassium's lilac).

Precipitates with sodium hydroxide and ammonia

• Adding sodium hydroxide solution to a solution of a metal compound identifies the positive ion (cation) it contains.
• Most metal hydroxides are insoluble in water. When mixed, the insoluble metal hydroxide forms as a solid suspended in the liquid - a precipitate.
• Method - Dissolve the sample in water, add the reagent a few drops at a time, record the precipitate colour, then add the reagent in excess. Test a fresh portion of the solution with aqueous ammonia.

Distinguishing the cations

• Aluminium vs zinc - Both dissolve in excess sodium hydroxide, but only zinc dissolves in excess ammonia.
• Copper(II) - Its blue precipitate dissolves in excess ammonia to give a distinctive deep blue solution.
• Calcium vs magnesium - Both give a white precipitate with sodium hydroxide that does not dissolve in excess. With ammonia, calcium gives little or no precipitate while magnesium gives a white one; a flame test also tells them apart.
• Always use a fresh sample portion for the ammonia test so leftover sodium hydroxide doesn't interfere.
• This test cannot identify Group 1 metals because their hydroxides are soluble.

Ionic equations for the hydroxide precipitates

An ionic equation shows only the reacting ions that form the precipitate, omitting unchanged spectator ions.

Constructing an ionic equation

1. Full balanced equation:
   • CuSO₄(aq) + 2NaOH(aq) → Cu(OH)₂(s) + Na₂SO₄(aq)
2. Split dissolved aqueous compounds into individual ions:
   • Cu²⁺(aq) + SO₄^2-(aq) + 2Na⁺(aq) + 2OH^-(aq) → Cu(OH)₂(s) + 2Na⁺(aq) + SO₄^2-(aq)
3. Cancel spectator ions (Na⁺ and SO₄^2-):
   • Cu²⁺(aq) + 2OH^-(aq) → Cu(OH)₂(s)

Common ionic equations

The number of OH^- ions always equals the charge on the metal cation so charges cancel to zero.

Worked example - Identifying a compound from tests

A white solid is thought to be a single ionic compound. Four tests are carried out. Name the compound and give its formula.

1. A flame test on the solid gives no colour.
2. Adding sodium hydroxide solution to a solution of the solid gives a green precipitate.
3. Adding dilute hydrochloric acid to a fresh portion of the solid produces no bubbles.
4. Acidifying a fresh portion of the solution with dilute hydrochloric acid, then adding barium chloride solution, gives a white precipitate.