Ion Identification

Identifying ions in solution

An ionic compound dissolved in water separates into mobile positive ions, called cations, and negative ions, called anions. Many of these solutions look alike, so appearance alone is usually not enough to identify an ion. Instead, a suitable reagent is added and the resulting observation is compared with known results.

These are qualitative tests: they identify which ion is present. They do not determine its concentration. In practical work, place small, separate portions of an unknown solution in clean test tubes. Using a fresh portion for each test prevents a reagent added in one test from affecting another.

Testing metal cations with sodium hydroxide

Add a few drops of dilute aqueous sodium hydroxide to the sample and shake gently. Hydroxide ions, $\text{OH}^-$, react with many dissolved metal ions to form insoluble metal hydroxides. The solid appears throughout the liquid as a precipitate.

For a metal ion $\text{M}^{n+}$, the general change is:

$$\text{M}^{n+}(aq)+n\text{OH}^-(aq)\rightarrow\text{M(OH)}_n(s)$$

The colour of the precipitate provides evidence about the cation:

For example, copper(II) ions form solid copper(II) hydroxide:

$$\text{Cu}^{2+}(aq)+2\text{OH}^-(aq)\rightarrow\text{Cu(OH)}_2(s)$$

Calcium and zinc ions both initially give white precipitates, so the first observation does not distinguish them. Continue adding sodium hydroxide until it is in excess. Zinc hydroxide dissolves to give a colourless solution, whereas the calcium hydroxide precipitate does not dissolve. The copper and iron precipitates also remain in excess sodium hydroxide.

The slow browning of the iron(II) precipitate is caused by oxidation in air. Record the initial green colour as well as any later change.

Testing for carbonate ions

Add a few drops of dilute acid to a fresh portion of the sample. Carbonate ions react with hydrogen ions to produce carbon dioxide and water:

$$2\text{H}^+(aq)+\text{CO}_3^{2-}(aq)\rightarrow\text{H}_2\text{O}(l)+\text{CO}_2(g)$$

Effervescence—visible bubbling—suggests that carbonate ions are present, but another gas-producing reaction could also bubble. Transfer the gas to limewater. Carbon dioxide turns limewater milky or cloudy white, confirming the result.

Testing for chloride, bromide and iodide ions

Chloride, bromide and iodide are halide ions. First acidify a fresh portion of the sample with dilute nitric acid, then add dilute silver nitrate solution. Silver ions combine with halide ions to form insoluble silver halides:

$$\text{Ag}^+(aq)+\text{X}^-(aq)\rightarrow\text{AgX}(s)$$

Here, $\text{X}^-$ represents $\text{Cl}^-$, $\text{Br}^-$ or ${\text{I}}^{-}$. The colour identifies the halide:

Acidification removes interfering carbonate ions, which could otherwise produce a precipitate and give a misleading result. Nitric acid is used rather than hydrochloric acid because hydrochloric acid would introduce chloride ions and could itself produce a white silver chloride precipitate.

Testing for sulfate ions

Acidify a fresh portion of the sample, then add dilute barium chloride solution or dilute barium nitrate solution. If sulfate ions are present, a white precipitate of barium sulfate forms:

$$\text{Ba}^{2+}(aq)+\text{SO}_4^{2-}(aq)\rightarrow\text{BaSO}_4(s)$$

A standard combination is dilute hydrochloric acid followed by barium chloride. Alternatively, dilute nitric acid and barium nitrate can be used. Acidifying first removes carbonate ions, because carbonate ions could also react with barium ions to form a white precipitate.

Drawing a conclusion from the evidence

An ion is identified by matching the whole test result: the correct reagent, the way it is used and the expected observation. For example, an unknown that gives a green precipitate with sodium hydroxide contains iron(II) ions, while an acidified unknown that gives a cream precipitate with silver nitrate contains bromide ions.

Precipitation is a chemical reaction, not an increase in mass. Ions that were already present and dispersed through the solution have simply joined to form an insoluble solid. In a closed system, total mass remains conserved.