Metal hydroxides — AQA GCSE Chemistry
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Metal hydroxides explained
Metal hydroxides covers 5 specification points, including Sodium hydroxide solution can be used to… and Solutions of aluminium, calcium and magnesium….
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Open a point below for its explanation and revision guidance.
Sodium hydroxide solution can be used to identify some metal ions (cations).
Adding sodium hydroxide solution to a solution containing metal cations produces insoluble metal hydroxides, which appear as coloured precipitates. The colour identifies the cation: for example, Cu²⁺ gives a blue precipitate, Fe²⁺ gives a green precipitate and Fe³⁺ gives a brown precipitate. For Al³⁺, Ca²⁺ and Mg²⁺ the precipitate is white, so further testing with excess sodium hydroxide is needed to tell aluminium apart. The method is to add a few drops of sodium hydroxide solution to the unknown solution, shake, and record the colour of any precipitate. This test identifies the cation, not the anion. The hydroxide ions combine with the metal cation to form the precipitate, while the sodium ions act as spectator ions and remain in solution.
Solutions of aluminium, calcium and magnesium ions form white precipitates when sodium hydroxide solution is added but only the aluminium hydroxide precipitate dissolves in excess sodium hydroxide solution.
Al³⁺, Ca²⁺ and Mg²⁺ all give a white precipitate with sodium hydroxide solution, so the initial observation cannot distinguish them. Adding excess sodium hydroxide solution resolves this: aluminium hydroxide dissolves to give a colourless solution, while calcium hydroxide and magnesium hydroxide remain as white precipitates. The method is to add sodium hydroxide solution dropwise, note the white precipitate, then continue adding excess and observe whether the precipitate disappears. If it dissolves, the cation is Al³⁺; if it remains, the cation is Ca²⁺ or Mg²⁺. This is a sequential test, and both observations should be reported.
Solutions of copper(II), iron(II) and iron(III) ions form coloured precipitates when sodium hydroxide solution is added.
Transition metal ions in aqueous solution can be identified by adding a few drops of dilute sodium hydroxide solution. The hydroxide ions react with the metal aqua ions to form an insoluble metal hydroxide, which appears as a coloured precipitate because the solid particles are dispersed through the liquid. Copper(II) ions give a blue precipitate, iron(II) ions give a green precipitate and iron(III) ions give a brown precipitate. For example, adding NaOH(aq) to copper(II) sulfate solution produces Cu(OH)₂(s), which settles as a blue solid. The colour is characteristic of the metal ion, so this test distinguishes the three ions from each other and from ions that give no precipitate or a white precipitate.
Copper(II) forms a blue precipitate, iron(II) a green precipitate and iron(III) a brown precipitate.
This statement gives the specific colour observed for each transition metal hydroxide precipitate. When sodium hydroxide solution is added to a solution containing copper(II) ions, a blue precipitate of copper(II) hydroxide forms. With iron(II) ions, the precipitate of iron(II) hydroxide is green. With iron(III) ions, the precipitate of iron(III) hydroxide is brown. These colours are distinctive and allow the ions to be identified in a simple test. For example, a green precipitate shows that the original solution contained iron(II) ions, not iron(III) ions, because the two oxidation states give different colours. The colour is an observation of the solid formed, not of the solution before reaction.
Students should be able to write balanced equations for the reactions to produce the insoluble hydroxides.
A balanced equation shows the reactants and products in the correct ratios and with the same number of each type of atom on both sides. For the precipitation reactions, the metal ion and hydroxide ion combine to form the insoluble metal hydroxide. For example, copper(II) sulfate solution reacts with sodium hydroxide solution to form copper(II) hydroxide precipitate and sodium sulfate solution: CuSO₄(aq) + 2NaOH(aq) → Cu(OH)₂(s) + Na₂SO₄(aq). The state symbols show that the hydroxide is a solid precipitate while the other compounds remain aqueous. Similar equations can be written for iron(II) and iron(III) hydroxides, taking care to balance the number of hydroxide ions against the charge on the metal ion.
Your focus
- Describe how sodium hydroxide solution is used to test for metal cations.
- Match precipitate colours to the correct metal ions.
- Explain that the precipitate is an insoluble metal hydroxide formed by the reaction of metal cations with hydroxide ions.
Show all 15 objectives
- Describe the white precipitates formed by Al³⁺, Ca²⁺ and Mg²⁺ with sodium hydroxide solution.
- Explain how excess sodium hydroxide solution distinguishes Al³⁺ from Ca²⁺ and Mg²⁺.
- Interpret observations from both stages of the test to identify the cation.
- Describe the appearance of the precipitates formed when sodium hydroxide solution is added to copper(II), iron(II) and iron(III) solutions.
- Match each coloured precipitate to the correct metal ion.
- Explain that the precipitate forms because the metal hydroxide produced is insoluble.
- Recall the colour of the precipitate formed by copper(II), iron(II) and iron(III) ions with sodium hydroxide solution.
- Use precipitate colour to identify an unknown transition metal ion.
- Distinguish between iron(II) and iron(III) ions from their precipitate colours.
- Write the correct formula for a metal hydroxide from the charge on the metal ion.
- Balance equations for the reaction of metal ion solutions with sodium hydroxide solution.
- Use state symbols correctly to show the insoluble hydroxide as a solid.
Metal hydroxides exam tips
Marking Points
- Add sodium hydroxide solution to the unknown solution and observe any precipitate.
- Record the colour of the precipitate, such as blue for Cu²⁺, green for Fe²⁺ or brown for Fe³⁺.
- Identify the cation from the precipitate colour, using the specified ions.
- Recognise that a white precipitate indicates Al³⁺, Ca²⁺ or Mg²⁺ and needs excess sodium hydroxide for further distinction.
- State that the precipitate is a metal hydroxide formed from the metal cation and hydroxide ions, with sodium ions remaining in solution.
- State that Al³⁺, Ca²⁺ and Mg²⁺ each give a white precipitate with sodium hydroxide solution.
- Add excess sodium hydroxide solution after the initial precipitate forms.
- State that aluminium hydroxide dissolves in excess sodium hydroxide solution, giving a colourless solution.
- State that calcium hydroxide and magnesium hydroxide do not dissolve in excess sodium hydroxide solution.
- Use the excess-sodium-hydroxide result to distinguish Al³⁺ from Ca²⁺ and Mg²⁺.
- Sodium hydroxide solution supplies hydroxide ions that react with metal ions in solution.
- The product is an insoluble metal hydroxide, so it appears as a precipitate rather than remaining dissolved.
- The colour of the precipitate depends on the identity of the metal ion present.
- Copper(II) ions produce a blue precipitate, iron(II) ions produce a green precipitate and iron(III) ions produce a brown precipitate.
- The test can be used to distinguish between solutions containing different transition metal ions.
- Copper(II) ions produce a blue precipitate with sodium hydroxide solution.
- Iron(II) ions produce a green precipitate with sodium hydroxide solution.
- Iron(III) ions produce a brown precipitate with sodium hydroxide solution.
- The colour identifies the metal ion and distinguishes between iron(II) and iron(III).
- The precipitate is the insoluble metal hydroxide, such as Cu(OH)₂, Fe(OH)₂ or Fe(OH)₃.
- Write the correct formulae for the reactants, including the metal compound and sodium hydroxide.
- Write the correct formula for the insoluble metal hydroxide product, using the metal ion charge to determine the number of hydroxide ions.
- Balance the equation so that the number of each type of atom is the same on both sides.
- Include state symbols where required, showing the metal hydroxide as a solid and the soluble reactants and by-product as aqueous.
- Check that the charges balance, for example Fe³⁺ requires three OH⁻ ions to form Fe(OH)₃.
Examiner Tips
- 💡Write the observation as colour plus precipitate, for example a blue precipitate, rather than colour alone.
- 💡If the precipitate is white, say that further testing with excess sodium hydroxide is needed before naming the cation.
- 💡Link each colour to the ion symbol with its charge, such as Cu²⁺ or Fe³⁺, to make the identification unambiguous.
- 💡Record both observations: the white precipitate first, then whether it dissolves in excess.
- 💡Use the phrase in excess sodium hydroxide solution to show the second stage has been carried out.
- 💡If the precipitate remains, name the possible ions as Ca²⁺ or Mg²⁺ and state that a further test is needed to separate them.
- 💡State the reagent as sodium hydroxide solution and the observation as a coloured precipitate, naming the colour precisely.
- 💡Link each colour to the correct ion, including the oxidation state in the name, for example iron(II) for green and iron(III) for brown.
- 💡If asked to identify an unknown, describe adding sodium hydroxide solution and observing the precipitate colour before naming the ion.
- 💡Learn the three colour–ion pairs as a table: copper(II) blue, iron(II) green, iron(III) brown.
- 💡In an identification question, give both the reagent and the colour, for example sodium hydroxide solution gives a green precipitate.
- 💡Use the oxidation state in the ion name so the examiner can see which iron ion is meant.
- 💡Work out the formula of the hydroxide first from the metal ion charge, then balance the rest of the equation around it.
- 💡Check each element in turn after balancing, including polyatomic ions such as sulfate, which can be counted as groups.
- 💡Include state symbols if the question asks for them, and make sure the precipitate is marked as solid.
Common Mistakes
- Adding excess sodium hydroxide immediately and missing the initial precipitate colour; correct this by adding a few drops first and observing before adding more.
- Confusing the green Fe²⁺ precipitate with the brown Fe³⁺ precipitate; correct this by learning the two colours as a pair.
- Stating that sodium ions form the precipitate; correct this by stating that hydroxide ions react with the metal cation to form the insoluble hydroxide.
- Stopping after the white precipitate and naming the ion; correct this by adding excess sodium hydroxide before deciding.
- Saying the aluminium precipitate disappears because it is washed away; correct this by stating that it dissolves in excess sodium hydroxide solution.
- Claiming calcium or magnesium hydroxide also dissolves in excess; correct this by stating that both remain as white precipitates.
- Describing the colour as belonging to the solution rather than the precipitate; the colour change is caused by the insoluble solid formed, so the observation should refer to the precipitate.
- Confusing iron(II) and iron(III) colours; iron(II) hydroxide is green while iron(III) hydroxide is brown, so the oxidation state must be matched to the correct colour.
- Writing that sodium hydroxide dissolves in the metal solution without reaction; the hydroxide ions react to form a new insoluble compound, so the word reaction or precipitate should be used.
- Swapping the colours of iron(II) and iron(III); the correct pairing is iron(II) green and iron(III) brown.
- Calling the copper precipitate green or the iron precipitate blue; copper(II) hydroxide is blue and should not be confused with iron colours.
- Forgetting that the colour is seen in the precipitate, not in the original solution; the observation should describe the solid formed after adding sodium hydroxide.
- Using the wrong number of hydroxide ions for the metal charge; for example, iron(III) needs three hydroxide ions, so the formula is Fe(OH)₃, not FeOH or Fe(OH)₂.
- Forgetting to balance the sodium and sulfate ions as well as the metal and hydroxide ions; the full equation must balance all atoms.
- Omitting state symbols or using the wrong state for the precipitate; the insoluble hydroxide should be shown as a solid.