A checklist of the mistakes that most often lose marks in the mole-concept chapter, grouped by content standard, for each one, what students write, why it loses marks, and the correct version to write instead.
Most marks in this chapter are lost not through misunderstanding the chemistry but through predictable slips in the calculation. Below are the errors examiners see again and again. For each, we show what students write, why it loses marks, and the correct version. Work through them with your own past attempts open beside you.
Relative mass (3.1)
1. Forgetting to multiply an Ar by the number of atoms. Students write Mr of H2O as 1 + 16 = 17. This loses the mark because there are two hydrogen atoms: the correct working is 2(1) + 16 = 18. Always read the subscript before adding.
2. Ignoring water of crystallisation. For CuSO4·5H2O students give 160 and stop. The five water molecules are part of the formula, so the correct Mr is 64 + 32 + 4(16) + 5(18) = 250. The dot in the formula does not mean “ignore”.
3. Miscounting atoms inside brackets. For Ca(OH)2 students write 40 + 16 + 1 = 57, applying the 2 to nothing. The bracket subscript multiplies everything inside: 40 + 2(16 + 1) = 74.
4. Giving Ar or Mr a unit. Writing “Mr of CO2 = 44 g” loses credit because relative mass is a ratio and has no unit. The unit g mol−1 belongs to molar mass, not to Mr.
The mole, particles and mass (3.2, 3.3)
5. Inverting the mole formula. Students compute moles as molar mass ÷ mass. The correct relationship is number of moles = mass ÷ molar mass. Writing the triangle out, n, m and M, before substituting prevents this.
6. Dividing mass by the Avogadro constant to get moles. To find moles you divide mass by molar mass, not by 6.02 x 1023 mol−1. The Avogadro constant converts between moles and number of particles, not between mass and moles.
7. Confusing molecules with atoms. Asked for the number of atoms in 1 mol of CO2, students give 6.02 x 1023 mol−1. That is the number of molecules; since each molecule has three atoms, the number of atoms is three times larger. Read whether the question asks for molecules, atoms or ions.
8. Forgetting units in the final answer. A correct number written as “0.5” instead of “0.5 mol”, or “24” instead of “24 dm3”, cannot earn the final mark. Every quantity needs its unit.
Volume of gas (3.4)
9. Using the wrong molar volume. Students apply 22.4 dm3 mol−1 when the question states room conditions, where the molar volume is 24 dm3 mol−1. Read the stated conditions and pick the matching value.
10. Mixing cm3 and dm3. A volume given as 500 cm3 is divided directly by a molar volume in dm3 mol−1. Convert first: 500 cm3 = 0.5 dm3, because 1 dm3 = 1000 cm3. Unit mismatch is one of the most common slips here.
11. Multiplying when you should divide (and the reverse). To get moles from volume you divide by the molar volume; to get volume from moles you multiply. Students often do the opposite. Keep the relationship n = V ÷ molar volume in front of you.
Empirical and molecular formula (3.5)
12. Not dividing by the smallest ratio. After finding moles of each element (say 0.2 and 0.6), students write the formula straight from those numbers. Divide both by the smallest (0.2) to get the simplest whole-number ratio, 1 : 3.
13. Rounding a ratio like 1.33 down to 1. A ratio of 1 : 1.33 is not 1 : 1; multiply both by 3 to get 3 : 4. Rounding a non-whole ratio changes the formula and loses the mark.
14. Giving the empirical formula when the molecular formula is asked. Students stop at CH2O when the question, having given the relative molecular mass, wants the molecular formula. Use n = Mr ÷ empirical formula mass and multiply every subscript by n.
Chemical equations and stoichiometry (3.6)
15. Using an unbalanced equation. Reading the mole ratio from H2 + O2 → H2O gives the wrong answer; the balanced equation 2H2 + O2 → 2H2O gives the ratio 2 : 1 : 2. Always balance before reading a ratio.
16. Using the mole ratio upside down. For 2Mg + O2 → 2MgO with 0.2 mol of oxygen, moles of MgO is 0.2 × (2 ÷ 1) = 0.4 mol, not 0.2 × (1 ÷ 2). Write the ratio in the direction “what you want over what you have”.
17. Treating coefficients or masses as moles. The 2 in 2Mg is a mole ratio, not 2 grams and not 2 atoms of mass. And you cannot use masses in a ratio directly, convert mass to moles first, apply the ratio, then convert back.
18. Omitting state symbols in Paper 2. When a question asks for a balanced equation, leaving out (s), (l), (g) or (aq) can cost a mark. Add the state symbols as a matter of habit.
19. Rounding too early in a multi-step calculation. Students round 3.33 to 3 partway through, so a ratio that should be 1 : 2 : 1 comes out wrong. Keep intermediate values to enough figures and round only at the very end, when you state the final answer with its unit.
20. Counting moles of molecules when moles of atoms are needed. In 0.5 mol of O2 there are 0.5 mol of oxygen molecules but 1.0 mol of oxygen atoms, because each molecule holds two atoms. Check whether the question counts molecules or atoms before you answer, especially when it then asks for a number of particles.
How to use this list
Turn each mistake into a check you run on your own work. Before you write a final answer, ask: is the equation balanced, did I convert to moles, is the ratio the right way up, and does every value carry a unit? A one-to-one teacher can mark a set of your calculations against exactly these checks and show you which slip is costing you most, so you stop repeating it. Because the same conversions run through acids and bases, electrochemistry and thermochemistry in later chapters of SPM Chemistry, fixing these habits now protects marks across the whole course. Keep this list beside the worked examples and the practice questions, and re-read it the night before any test on the mole concept.
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