What you'll learn
This revision guide covers the analytical techniques required to identify metal ions in unknown compounds, a core practical skill in AQA GCSE Chemistry. You'll learn how to conduct flame tests safely and accurately, interpret characteristic flame colours, and use additional chemical tests to confirm the identity of specific metal cations. These techniques are essential for both written examinations and required practicals.
Key terms and definitions
Flame test — a qualitative analytical technique where a sample is heated in a Bunsen burner flame to observe characteristic colours produced by metal ions
Cation — a positively charged ion formed when an atom loses one or more electrons; metal ions are cations
Precipitate — an insoluble solid formed when two solutions react together
Sodium hydroxide solution — an alkali (NaOH) used to test for the presence of metal ions by forming characteristic precipitates
Qualitative analysis — analytical methods that identify which substances are present, rather than measuring how much
Spectroscopy — modern instrumental methods of analysis that can identify elements and compounds, including flame emission spectroscopy and atomic absorption spectroscopy
Lithium compounds — ionic compounds containing Li⁺ ions that produce a crimson flame colour
Transition metals — elements between Groups 2 and 3 in the periodic table that form coloured compounds and ions
Core concepts
Why flame tests work
When metal compounds are heated strongly in a flame, the energy excites electrons within the metal ions. These electrons move from their normal energy levels (ground state) to higher energy levels. As the electrons fall back down to their original positions, they release this absorbed energy as light of specific wavelengths. Different metals release different amounts of energy, producing characteristic colours that we can observe.
This emission of light at specific wavelengths is unique to each element, making flame tests a reliable identification method. The technique works best with Group 1 and Group 2 metals, which produce particularly distinctive flame colours.
Conducting a flame test safely
The correct procedure is essential for accurate results and safety:
- Dip a clean nichrome or platinum wire loop into concentrated hydrochloric acid to clean it
- Heat the wire in the blue Bunsen flame until no colour is seen (this removes any previous contamination)
- Dip the clean wire into the hydrochloric acid again
- Dip the wire into the solid sample or solution being tested
- Place the wire and sample into the edge of the blue Bunsen flame
- Observe and record the flame colour produced
- Clean the wire thoroughly between tests to prevent cross-contamination
Safety considerations:
- Wear eye protection throughout
- Ensure the Bunsen burner is on a heatproof mat
- Keep long hair tied back and loose clothing away from flames
- Use concentrated hydrochloric acid carefully as it is corrosive
- Do not look directly at very bright flames
Flame colours for common metal ions
You must memorise these characteristic flame colours for the AQA examination:
Lithium (Li⁺) — crimson or red flame
Sodium (Na⁺) — yellow or orange flame (often persistent and very bright)
Potassium (K⁺) — lilac flame (pale purple, often difficult to see)
Calcium (Ca²⁺) — orange-red flame (sometimes described as brick red)
Copper (Cu²⁺) — green or blue-green flame
Barium (Ba²⁺) — pale green flame
Note that sodium contamination is very common in laboratories because sodium compounds are everywhere. Even traces of sodium will produce a strong yellow colour that can mask other flame colours. This is why cleaning the wire thoroughly between tests is critical.
Using sodium hydroxide solution to identify metal ions
Not all metal ions produce distinctive flame colours. Sodium hydroxide solution provides an alternative test that forms coloured precipitates with many metal ions. When sodium hydroxide solution is added dropwise to a solution containing metal ions, characteristic precipitates form:
Calcium (Ca²⁺) — white precipitate (if enough is present; often barely visible)
Copper (Cu²⁺) — blue precipitate of copper(II) hydroxide
Iron(II) (Fe²⁺) — green precipitate of iron(II) hydroxide (may turn brown on standing as it oxidises to iron(III))
Iron(III) (Fe³⁺) — brown or rust-coloured precipitate of iron(III) hydroxide
Aluminium (Al³⁺) — white precipitate initially; dissolves in excess sodium hydroxide to form a colourless solution
Magnesium (Mg²⁺) — white precipitate of magnesium hydroxide
The chemical reaction occurring is a precipitation reaction. For example, with copper ions:
Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s)
Distinguishing between similar precipitates
Several metal ions form white precipitates with sodium hydroxide, making them difficult to distinguish. Further testing helps identify them:
Calcium produces a weak white precipitate but can be confirmed by its distinctive orange-red flame test.
Magnesium produces a white precipitate and shows no distinctive flame colour.
Aluminium initially forms a white precipitate, but this is the key distinguishing feature: the precipitate dissolves when excess sodium hydroxide is added. The aluminium hydroxide is amphoteric, meaning it can react with both acids and alkalis.
Limitations and modern methods
Flame tests have important limitations:
- Some metal ions produce no distinctive colour
- Very similar colours can be confused (e.g., calcium and sodium both produce orange-type colours)
- Mixtures of metals are difficult to identify
- The method requires experience to interpret colours accurately
- Contamination (especially by sodium) causes false results
Instrumental methods used professionally include:
Flame emission spectroscopy — measures the specific wavelengths of light emitted by excited metal ions with much greater accuracy than visual observation. The intensity of light at particular wavelengths indicates which elements are present and their concentrations.
Atomic absorption spectroscopy — measures how much light of specific wavelengths is absorbed by a sample, identifying elements present and their quantities.
These modern techniques offer advantages over simple flame tests:
- Much more sensitive (can detect trace amounts)
- More accurate and precise
- Can analyse mixtures effectively
- Provide quantitative data (how much, not just what)
- Less affected by contamination
- Results can be stored and compared electronically
However, simple flame tests remain valuable because they require minimal equipment, give instant results, and work well for common metal ions in relatively pure samples.
Worked examples
Example 1: Identifying an unknown metal compound
Question: A student performs a flame test on an unknown white powder. The powder produces a crimson flame. The student then dissolves some powder in water and adds sodium hydroxide solution dropwise. No precipitate forms. Identify the metal ion present and explain your reasoning. (3 marks)
Answer:
- The metal ion is lithium (Li⁺) [1 mark]
- The crimson flame colour is characteristic of lithium ions [1 mark]
- No precipitate forms because lithium hydroxide is soluble in water [1 mark]
Examiner comment: This question tests both flame test knowledge and understanding that not all metal hydroxides are insoluble. Group 1 metal hydroxides (lithium, sodium, potassium) are all soluble, so they don't form precipitates with sodium hydroxide solution.
Example 2: Distinguishing between two metal ions
Question: A compound contains either calcium ions or magnesium ions. Describe two tests you could perform to distinguish between these ions, including the expected observations. (4 marks)
Answer:
Test 1: Add sodium hydroxide solution to a solution of the compound
- Both form white precipitates [1 mark]
- So this test alone cannot distinguish them
Test 2: Perform a flame test on the solid compound
- Calcium produces an orange-red flame [1 mark]
- Magnesium produces no distinctive flame colour / very faint white [1 mark]
- This confirms which ion is present [1 mark]
Examiner comment: Questions asking you to distinguish between ions are common. You must state what you do and what you observe for each ion. Simply naming tests without observations won't earn full marks.
Example 3: Analysing hydroxide precipitate reactions
Question: A solution contains iron(III) ions. When sodium hydroxide solution is added, a brown precipitate forms.
(a) Write the ionic equation for this reaction. (2 marks) (b) Suggest why this test is more reliable than a flame test for identifying iron ions. (1 mark)
Answer:
(a) Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s) [2 marks: 1 for correct formula and charges, 1 for balanced equation with state symbols]
(b) Iron ions do not produce a distinctive flame colour / flame tests don't work well for transition metals / the brown precipitate is very characteristic and easy to identify [1 mark]
Examiner comment: For ionic equations at GCSE, ensure you include: correct charges, balanced equation, and state symbols (aq) for aqueous and (s) for solid precipitate. Transition metals generally don't produce good flame colours, so the hydroxide precipitation test is more reliable.
Common mistakes and how to avoid them
Confusing flame colours — Calcium (orange-red) and sodium (yellow-orange) can look similar. Remember sodium flames are pure yellow and very bright, while calcium has more red in the colour. Potassium lilac is faint and easily missed. Practice describing colours precisely using the correct terminology.
Not cleaning the wire properly — Failing to clean the wire between tests leads to contamination, especially from sodium. Always clean until no colour appears in the flame. The wire must be dipped in acid and heated to colourless between each test.
Stating that precipitates "disappear" in excess sodium hydroxide — Use correct chemical terminology: aluminium hydroxide "dissolves" in excess sodium hydroxide, forming a soluble compound. Precipitates don't vanish; they undergo further chemical reactions.
Forgetting that Group 1 hydroxides are soluble — Sodium, potassium, and lithium hydroxides don't form precipitates with sodium hydroxide solution. Don't expect to see a precipitate for every metal ion.
Incomplete descriptions in exam answers — When asked to describe a test, you must state: (1) what you add/do, (2) what you observe, and (3) what this indicates. Missing any element loses marks.
Mixing up iron(II) and iron(III) precipitate colours — Iron(II) gives a green precipitate (may turn brown on standing), while iron(III) gives a brown precipitate immediately. The Roman numerals indicate different oxidation states of iron.
Exam technique for chemical analysis questions
Command word "Describe" — You must state both the method and the observations. For flame tests: "Dip a clean wire in the sample, hold in a flame, observe [colour] flame." For precipitate tests: "Add sodium hydroxide solution dropwise, observe [colour] precipitate forms."
Writing ionic equations — Only include the ions that react (the precipitate-forming ions). Sodium and other spectator ions are omitted. Always include state symbols: (aq) for aqueous, (s) for solid precipitate. Check charges balance on both sides.
Justifying your answer — When identifying an unknown, link observations to ions: "The lilac flame indicates potassium ions because this colour is characteristic of K⁺." Simply stating "potassium" without explanation loses marks.
Practical questions — Marks are often awarded for: safety precautions, accurate method, observations, and conclusions. Required practical questions may ask you to evaluate results or suggest improvements. Always relate improvements to specific issues (contamination, accuracy, safety).
Quick revision summary
Flame tests identify metal ions by their characteristic colours when heated: lithium produces crimson, sodium gives yellow, potassium shows lilac, calcium creates orange-red, copper produces green, and barium gives pale green flames. Clean the wire thoroughly between tests to prevent sodium contamination. Sodium hydroxide solution provides an alternative test, forming coloured precipitates: blue for copper(II), green for iron(II), brown for iron(III), and white for calcium, magnesium, and aluminium. Aluminium hydroxide uniquely dissolves in excess sodium hydroxide. Modern instrumental methods like flame emission spectroscopy offer greater sensitivity and accuracy but simple chemical tests remain useful for rapid identification.