PAG4: Qualitative Analysis of Ions

Plan tests for the candidate cations and anions in the task. The examples here include Group 2, ammonium, halides, carbonate, hydroxide and sulfate, followed by transition-ion comparisons. Use labelled fresh portions so earlier reagents cannot contaminate later tests. Record the reagent, conditions and observation before inferring an identity.

For the familiar Group 2 flame tests, calcium gives brick-red/orange-red, strontium crimson and barium apple-green; magnesium has no characteristic visible flame colour in this test. Sodium contamination can mask weaker colours. Follow the approved clean-wire procedure and do not infer that every Group 2 ion can be identified reliably from a visible flame alone. Solubility trends can provide supporting comparisons.

Carbonate releases carbon dioxide on addition of dilute acid; passing the gas through limewater gives cloudiness. Hydroxide is alkaline, so an indicator or pH measurement supports its detection, but alkalinity alone cannot distinguish hydroxide from every other basic anion. Interpret the set of possible unknowns and use complementary tests.

For sulfate, acidify with dilute hydrochloric acid then add barium chloride: a white barium sulfate precipitate persists. Acid removes carbonate interference; sulfuric acid would introduce sulfate. For halides use dilute nitric acid then silver nitrate: chloride gives white, bromide cream and iodide yellow precipitates. Hydrochloric acid would introduce chloride. Silver chloride dissolves in dilute ammonia, silver bromide in concentrated ammonia, and silver iodide remains insoluble under those tests.

With aqueous sodium hydroxide and gentle warming, ammonium ions release ammonia, which turns damp red litmus blue. Use the teacher’s ventilation and warming arrangements; do not identify a gas by directly inhaling it. The ionic equation is NH₄⁺ + OH⁻ → NH₃ + H₂O.

Small test portions reduce reagent use and hazardous waste. Barium and silver salts and ammonia require the specified handling and disposal. A matching observation supports an identification within the candidate set; mixtures, low concentration or contamination can create ambiguity. Repeat with fresh apparatus or compare a known reference when the observation is weak rather than inventing a more decisive colour.

Compare the metal ions specified in the task, including iron(II), iron(III) and copper(II) in these examples. Aluminium provides a useful comparison but is not itself a transition metal. Keep sample concentration, reagent addition and observation conditions consistent.

Add hydroxide dropwise, record the initial precipitate, then add the specified excess and record whether it dissolves. Copper(II) normally gives a blue precipitate, iron(II) green and iron(III) brown. Iron(II) hydroxide can brown on standing as it oxidises in air, so the time of observation matters. Aluminium gives a white precipitate that dissolves in excess hydroxide through amphoteric behaviour.

Ammonia is a weak base and can remove protons from water ligands, causing hydroxide precipitation. In excess it can also act as a ligand. With copper(II), the initial blue precipitate dissolves in sufficient excess ammonia to form a deep-blue solution of the ammine complex. This is ligand substitution, not reduction of copper(II) to copper metal.

For copper, the simplified precipitation equation is Cu²⁺ + 2OH⁻ → Cu(OH)₂(s). An aqua-complex representation shows the water ligands and proton removal more explicitly. For excess ammonia, [Cu(H₂O)₆]²⁺ + 4NH₃ ⇌ [Cu(NH₃)₄(H₂O)₂]²⁺ + 4H₂O. Check atoms and charge. Iron hydroxides do not show the same dissolution in excess ammonia under the usual test conditions.

Use fresh separate portions for hydroxide, ammonia and any carbonate comparison required by the specification. A mixed unknown or coloured starting solution can obscure a precipitate. A known reference, clean background and fresh iron(II) solution can improve interpretation. Do not report an initial “blue solution” as a “blue precipitate”: physical state is part of the evidence.

Ammonia vapour and metal-ion solutions require the specified ventilation, eye protection and waste collection. Adding a large uncontrolled excess from the start can conceal the initial precipitate. A useful improvement is dropwise addition with a recorded sequence, followed by enough excess to test solubility. Repeating a contaminated test with the same tube will not resolve the source of the error.

Attempt these problems on paper before opening the worked reasoning. Explain your steps as well as giving a final answer.