The relationship between empirical formula and molecular formula.

The idea

The molecular formula is always a whole-number multiple of the empirical formula, so molecular = n × empirical where n = relative molecular mass ÷ empirical formula mass. Work out the empirical formula mass by adding the relative atomic masses in the empirical unit, divide the relative molecular mass by it, then multiply every subscript in the empirical formula by that value. For an empirical formula of CH₂O with mass 30.0 and a relative molecular mass of 180.0, n is 6 and the molecular formula is C₆H₁₂O₆. The multiplier must be a whole number: a value such as 2.9 signals an arithmetic error or a rounded relative molecular mass rather than a fractional formula.

What earns the marks

• The empirical formula mass obtained by summing the relative atomic masses in the empirical unit.
• Dividing the relative molecular mass by the empirical formula mass to find the multiplier.
• Multiplying every subscript in the empirical formula by that multiplier.
• The multiplier given as a whole number.

Common mistakes and how to put them right

• Multiplying only the first element's subscript by the multiplier.

• Dividing the empirical formula mass by the relative molecular mass, inverting the ratio.

• Rounding a multiplier of 2.9 down to 2 rather than checking the earlier working.

• Forgetting that the multiplier can legitimately be 1.

Exam tips

• Write the multiplier down as a labelled value before you use it; it is worth a mark on its own.
• Check the finished molecular formula by recalculating its relative molecular mass against the value given.
• If the multiplier is not close to a whole number, recheck the empirical formula rather than forcing the answer.