A student analyses a compound containing only iron and sulfur. The student finds that 2.24 g of iron combines with 1.92 g of sulfur to form the compound. Use the relative atomic masses Fe = 56, S = 32 to determine the empirical formula of the compound. Show your reasoning at each step.

OCR A-Level Chemistry A (H432) — 2.2 Amount of substance · Explain · 5 marks · View as Markdown

Written & reviewed by James Millett — Biology (Imperial College London), PGCE Science (University of Cambridge).

Iron sulfide compounds are used in the manufacture of certain pigments. Determining the empirical formula of such compounds is an important analytical task in the chemical industry.

Model answer (5 marks)

1. Calculate moles of Fe: 2.24 g ÷ 56 g mol⁻¹ = 0.040 mol Fe
2. Calculate moles of S: 1.92 g ÷ 32 g mol⁻¹ = 0.060 mol S
3. Divide by the smaller number (0.040 mol): Fe = 0.040/0.040 = 1, S = 0.060/0.040 = 1.5
4. Multiply the ratio by 2 to give whole numbers: Fe = 2, S = 3
5. Empirical formula: Fe₂S₃

Examiner tips

  • Show each calculation step clearly; include units. Use the smallest mole value to normalise the ratio. Multiply by the smallest integer that gives whole numbers. State the final empirical formula explicitly.

Common mistakes

  • Using the wrong atomic mass (e.g., 55.8 instead of 56). Failing to normalise the mole ratio before multiplying. Giving the formula as FeS₁.₅ instead of Fe₂S₃.

Mark scheme (5 marks)

  1. Calculates the number of moles of iron: 2.24 ÷ 56 = 0.04 mol
  2. Calculates the number of moles of sulfur: 1.92 ÷ 32 = 0.06 mol
  3. Divides both values by the smallest molar amount (0.04) to find the ratio
  4. Recognises that a ratio of 1 : 1.5 must be multiplied by 2 to give whole numbers, giving Fe : S = 2 : 3
  5. States the correct empirical formula: Fe₂S₃

Key terms in this question

empirical formula · relative atomic mass

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