Identification of Cations and Anions

Principles and Mechanics of Flame Tests

• Purpose of Flame Tests: Metals change the color of a flame when heated in it, producing a distinct color that serves as a diagnostic tool for identifying common metals.
• Required Equipment: A platinum or nichrome wire is used for the procedure.
• Cleaning Procedure:     * The wire must be cleaned by dipping it into concentrated $HCl_{(aq)}$ (hydrochloric acid).     * The wire is then held in a hot Bunsen flame.     * This cleaning process (dipping and heating) is repeated until no color is produced in the Bunsen flame.
• Sample Preparation and Observation:     * The wire may be looped at the end or left straight.     * Before testing the sample, the wire is wetted with some of the concentrated acid.     * The wet wire is dipped into a small amount of the compound containing the ion to be identified.     * The wire is placed back into the flame, and the resulting color is observed.
• Theoretical Basis for Color Changes:     * Excitation: The distinct colors are produced when electrons in the metal ions absorb energy from the heat and transition (jump) to higher energy levels. This relates to the concept of the absorption spectrum.     * Relaxation: When these electrons fall back down to lower energy levels, they emit light energy in the form of specific wavelengths. These observed wavelengths constitute the emission spectrum.

Characteristic Flame Colors for Cations

• Alkali Metals (Group 1):     * $Li^+$ (Lithium): Red     * $Na^+$ (Sodium): Yellow-Orange     * $K^+$ (Potassium): Lilac (pink)     * $Rb^+$ (Rubidium): Red-Violet
• Alkaline Earth Metals (Group 2):     * $Mg^{2+}$ (Magnesium): White     * $Ca^{2+}$ (Calcium): Orange-Red     * $Sr^{2+}$ (Strontium): Red     * $Ba^{2+}$ (Barium): Pale Green
• Post-Transition and Transition Metals:     * $Pb^{2+}$ (Lead): Grey-White     * $Cu^{2+}$ (Copper): Blue-Green

Cation Identification Using Sodium Hydroxide ($NaOH$)

Sodium hydroxide solution is added to solutions containing metal ions to observe the color of the precipitate ($ppt$) formed and its solubility in excess base.

• Iron ($Fe^{2+}$):     * Observation: Pale green precipitate produced.     * Effect of Excess $NaOH$: Precipitate remains (insoluble).
• Iron ($Fe^{3+}$):     * Observation: Red-orange precipitate produced.     * Effect of Excess $NaOH$: Precipitate remains (insoluble).
• Copper ($Cu^{2+}$):     * Observation: Blue precipitate produced.     * Effect of Excess $NaOH$: Precipitate remains (insoluble).
• Manganese ($Mn^{2+}$):     * Observation: Pale brown precipitate produced.     * Effect of Excess $NaOH$: Precipitate remains (insoluble).
• Chromium ($Cr^{3+}$):     * Observation: Green precipitate produced.     * Effect of Excess $NaOH$: Precipitate dissolves (very soluble in excess).
• Magnesium ($Mg^{2+}$):     * Observation: White precipitate produced.     * Effect of Excess $NaOH$: Precipitate remains (insoluble).
• Aluminium ($Al^{3+}$):     * Observation: White precipitate produced.     * Effect of Excess $NaOH$: Precipitate dissolves.
• Calcium ($Ca^{2+}$):     * Observation: White precipitate produced.     * Effect of Excess $NaOH$: Precipitate remains (insoluble).
• Zinc ($Zn^{2+}$):     * Observation: White precipitate produced.     * Effect of Excess $NaOH$: Precipitate dissolves.
• Lead ($Pb^{2+}$):     * Observation: White precipitate produced.     * Effect of Excess $NaOH$: Precipitate dissolves.
• Barium ($Ba^{2+}$):     * Observation/Excess: No precipitate ($No\,ppt$).

Cation Identification Using Aqueous Ammonia ($NH_4OH$)

Aqueous ammonia solution ($NH_3$ in water) is added dropwise and then in excess to metal ion solutions.

• Iron ($Fe^{2+}$):     * Observation: Pale green precipitate produced.     * Effect of Excess $NH_4OH$: Precipitate remains.
• Iron ($Fe^{3+}$):     * Observation: Red-orange precipitate produced.     * Effect of Excess $NH_4OH$: Precipitate remains.
• Copper ($Cu^{2+}$):     * Observation: Blue precipitate produced.     * Effect of Excess $NH_4OH$: Precipitate dissolves to form a deep blue solution.
• Manganese ($Mn^{2+}$):     * Observation: Pale brown precipitate produced.     * Effect of Excess $NH_4OH$: Precipitate remains.
• Chromium ($Cr^{3+}$):     * Observation: Grey-Green precipitate produced.     * Effect of Excess $NH_4OH$: Precipitate remains.
• Magnesium ($Mg^{2+}$):     * Observation: White precipitate produced.     * Effect of Excess $NH_4OH$: Precipitate remains.
• Aluminium ($Al^{3+}$):     * Observation: White precipitate produced.     * Effect of Excess $NH_4OH$: Precipitate remains.
• Zinc ($Zn^{2+}$):     * Observation: White precipitate produced.     * Effect of Excess $NH_4OH$: Precipitate dissolves.
• Lead ($Pb^{2+}$):     * Observation: White precipitate produced.     * Effect of Excess $NH_4OH$: Precipitate remains.
• Calcium ($Ca^{2+}$) and Barium ($Ba^{2+}$):     * Observation/Excess: No precipitate ($No\,ppt$).

Tests for the Ammonium Ion ($NH_4^+$)

• Addition of Sodium Hydroxide ($NaOH$): No precipitate is formed initially upon the addition of sodium hydroxide.
• Heating Procedure: If the solution is heated following the addition of $NaOH$, ammonia gas ($NH_3$) will evolve.
• Identification of Ammonia Gas:     * Ammonia gas has a characteristic choking smell.     * Ammonia gas will turn moist (damp) red litmus paper blue.

Identification of the Nitrate Ion ($NO_3^-$)

• Concentrated Sulphuric Acid Test:     * Effervescence is observed during the reaction.     * The gas evolved will form a white precipitate when bubbled through lime water.
• Heat Test:     * A sample of the nitrate salt is heated directly.     * The evolved gas will form a white precipitate with lime water.
• Brown Ring Test:     1. Iron(II) sulphate ($FeSO_4$) solution is added to the nitrate solution.     2. Concentrated sulphuric acid ($H_2SO_4$) is added slowly down the side of the test tube so that it forms a separate layer below the aqueous solution.     3. A distinct brown ring forms at the junction where the two liquid layers meet.
• Devarda's Alloy Test:     * Devarda's alloy is a mixture containing Copper ($Cu$), Aluminium ($Al$), and Zinc ($Zn$).     1. The alloy mixture is added to the nitrate solution in the presence of sodium hydroxide ($NaOH$).     2. Ammonia ($NH_3$) gas is liberated.     3. The ammonia is confirmed by its ability to turn damp red litmus paper blue.
• Copper Turnings Test:     1. Copper turnings and sulphuric acid are added to the nitrate solution and then heated.     2. The copper dissolves into the solution.     3. The nitrate decomposes to produce nitrogen dioxide ($NO_2$) gas.     * Chemical Equation: $Cu_{(s)} + 4HNO_{3(aq)} \rightarrow Cu(NO_3)<em>{2(aq)} + 2NO</em>{2(g)} + 2H_2O_{(l)}$

General Solubility Rules for Salts

Soluble Salts

• Nitrates, Nitrites, and Acetates: All are soluble.
• Sulphates: All sulphates are soluble except for Barium sulphate ($BaSO_4$) and Lead sulphate ($PbSO_4$).     * Calcium sulphate ($CaSO_4$) and Silver sulphate ($Ag_2SO_4$) are categorized as sparingly soluble.
• Halides: Most halides (chlorides, bromides, iodides) are soluble except for those of Silver and Lead.     * Lead chloride ($PbCl_2$) and Lead bromide ($PbBr_2$) are soluble in hot water.
• Hydrogen Carbonates: Most hydrogen carbonates ($HCO_3^-$) are soluble.

Insoluble Salts

• Carbonates: All carbonates are insoluble except for those of Sodium ($Na^+$), Potassium ($K^+$), and Ammonium ($NH_4^+$).