Edexcel IGCSE Chemistry · Spec 1.31-1.36
Empirical & Molecular Formulae
How to work out empirical and molecular formulae from experimental data.
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Empirical & Molecular Formulae
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Explained
Two kinds of formula, and how to get from data to one of them
The molecular formula gives the actual number of atoms of each element in one molecule. The empirical formula gives the simplest whole number ratio of those atoms.
Glucose is C6H12O6 as a molecular formula. Divide every number by six and the ratio is one carbon to two hydrogens to one oxygen, so its empirical formula is CH2O. Ethene is C2H4, so its empirical formula is CH2, and propene is C3H6, which gives CH2 as well. Two different compounds can share an empirical formula, which is why the molecular one is needed to identify a substance.
Working out an empirical formula
The same four steps work every time, whether you are given masses or percentages.
- Write down the mass, or the percentage, of each element.
- Divide each by that element's relative atomic mass. This converts mass into moles, which is the number of atoms.
- Divide every answer by the smallest of them, which turns the numbers into a ratio.
- Round to whole numbers and write the formula.
If a number comes out close to something like 1.5 or 2.5, do not round it. Multiply every number by two instead, because a ratio of 1 to 1.5 is really 2 to 3.
Getting the molecular formula from it
You need the relative formula mass of the compound, which the question will give you.
Find the mass of one empirical formula unit, divide the relative formula mass by it, and multiply every subscript in the empirical formula by that answer.
CH2O has a mass of 30. If the compound's relative formula mass is 180, then 180 divided by 30 is 6, so the molecular formula is C6H12O6.
What the mark scheme accepts and rejects
An Edexcel International GCSE Chemistry mark scheme for a three mark question on propene asks for three different things in three boxes: the displayed or molecular formula, the empirical formula, and the relative formula mass. Beside the first it says the answer must show all the bonds, and beside the last it says to ignore the calculation.
Two lessons come out of that. A displayed formula has to show every bond, including the ones to hydrogen, and a partly drawn structure gets nothing. And where only the final value is being marked, working is not required, though writing it costs nothing and helps you check.
The examiner report for a paper in the same series notes that the empirical formula calculation was answered successfully by most learners, but that a value of 6.94 for hydrogen was occasionally rounded down to 6 rather than up to 7. It also records that relatively few learners used atomic numbers instead of relative atomic masses, or produced an inverted calculation, both of which had been common errors in previous years.
Dividing by the atomic number instead of the relative atomic mass is the error to watch for. Carbon's atomic number is 6 and its relative atomic mass is 12, and only one of those belongs in this calculation.
Finding the formula of an oxide by experiment
Weigh an empty crucible, add a known mass of magnesium ribbon, and weigh again. Heat it strongly with the lid slightly raised, so oxygen can get in but the white smoke of magnesium oxide cannot escape.
Heat, cool and weigh repeatedly until the mass stops changing, which shows the reaction is complete. The increase in mass is the mass of oxygen that has combined.
You now have the mass of magnesium and the mass of oxygen, so the four steps give you MgO. The two common sources of error are losing some of the product as smoke, which makes the oxygen mass too low, and not heating for long enough, which does the same.
Spec 1.31-1.36
What you need to know
- Tell the difference between an empirical and a molecular formula
- Work out an empirical formula from masses
- Find the formula of an oxide by experiment
Active recall
Quick check
Answer each question before opening the answer.
What is the difference between an empirical and a molecular formula?
Empirical formula is the simplest whole-number ratio of atoms; molecular formula is the actual number of each atom in a molecule.
How do you find an empirical formula from masses or percentages?
Divide each element's mass/percentage by its Ar, then divide all answers by the smallest to get the simplest ratio.
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