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1.5.8Principles of Chemistry

Empirical and Molecular Formulae — Edexcel International GCSE Science Double Award 4SD0 - Chemistry Revision Notes & Topical Questions

5 min read 12 parts Lesson 31 of 50

📚 Edexcel International GCSE Science Double Award 4SD0 - Chemistry Hub / Topical Revision
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Part 1 of 5Two kinds of formula

Empirical formula
  • Empirical formula This shows the simplest whole-number ratio of atoms of each element in a compound.
  • Ethane, C2H6, has the empirical formula CH3.
Molecular formula
  • Molecular formula This shows the actual number of atoms of each element in one molecule.
  • It is always a whole-number multiple of the empirical formula.
When they are the same
  • When they are the same Where the ratio cannot be simplified, as in H2O or CH4, the two formulae are identical.
    • Ionic compounds are given only as empirical formulae.

Part 2 of 5Finding the empirical formula

1
Write down the mass of each element, or take the percentages as masses in a 100 g sample.
2
Divide each mass by the relative atomic mass of the element to get moles.
3
Divide every mole value by the smallest one to get a ratio.
4
If any number is not close to a whole number, such as 1.5, multiply all of them by a small integer, here 2, to clear the fraction.
5
Write the formula using the whole numbers as subscripts.

Part 3 of 5Worked example: empirical formula

StepCarbonHydrogen
Data: percentage by mass80%20%
Mass in 100 g80 g20 g
Divide by Ar (C = 12, H = 1)80 ÷ 12 = 6.6720 ÷ 1 = 20
Divide by the smaller (6.67)13
Empirical formulaCH3

Part 4 of 5From empirical to molecular formula

1Find the empirical formula mass
  • Find the empirical formula mass Add up the Ar values in the empirical formula. For CH3 this is 12 + 3 = 15.
2Compare with the Mr
  • Compare with the Mr The question gives the relative formula mass of the molecule, for example 30.
3Find the multiple
  • Find the multiple Divide the molecular Mr by the empirical formula mass: 30 ÷ 15 = 2.
4Multiply the formula
  • Multiply the formula Multiply every subscript by 2 to get C2H6, the molecular formula.

Part 5 of 5Using combustion or reaction data

Masses of products
  • Masses of products When an organic compound burns, all the carbon ends up as CO2 and all the hydrogen as H2O,
    • so the masses of these products can be converted back to masses of C and H.
Oxygen by difference
  • Oxygen by difference If a compound contains oxygen too, its mass is found by subtracting the masses of the other elements from the total sample mass.
Rounding sensibly
  • Rounding sensibly Experimental values like 1.98 or 2.03 are rounded to 2,
    • but values such as 1.5 or 2.33 signal that you must multiply to make whole numbers.

Key terms

TermMeaning
Empirical formulaThe simplest whole-number ratio of the atoms of each element in a compound.
Molecular formulaThe actual number of atoms of each element in one molecule of a compound.
Relative atomic mass (Ar)The average mass of an atom of an element compared with one twelfth of a carbon-12 atom.
Relative formula mass (Mr)The sum of the relative atomic masses of all the atoms in the formula.
MoleThe amount of substance containing 6.02 × 1023 particles.

Avoid these mistakes

1
Dividing percentages by atomic numbers instead of relative atomic masses.
2
Rounding 1.5 to 2 instead of multiplying the whole ratio by 2.
3
Stopping at the empirical formula when the question asks for the molecular formula and gives an Mr.
4
Multiplying only one subscript by the multiple rather than all of them.

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