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S1.1 - Introduction to the particulate nature of matter S1.2 - The nuclear atom S1.3 - Electron configurations S1.4 - Counting particles by mass - The mole S1.5 - Ideal gases S2.1 - The ionic model S2.2 - The covalent model S2.3 - The metallic model S2.4 - From models to materials S3.1 - The periodic table - Classification of elements S3.2 - Functional groups - Classification of organic compounds R1.1 - Measuring enthalpy changes R1.2 - Energy cycles in reactions R1.3 - Energy from fuels R1.4 - Entropy and spontaneity AHL R2.1 - How much? The amount of chemical change R2.2 - How fast? The rate of chemical change R2.3 - How far? The extent of chemical change R3.1 - Proton transfer reactions R3.2 - Electron transfer reactions R3.3 - Electron sharing reactions R3.4 - Electron-pair sharing reactions

S1.4 - Counting particles by mass - The Mole

1.4.1 The Mole and Avagadro's Constant 1.4.2 Relative Atomic Mass (Ar) and Relative Formula Mass (Mr) 1.4.3 Molar Mass, Mass and the Mole 1.4.4 Empirical and Intermolecular Formulae 1.4.5 Molar Concentration and Solution Calculations 1.4.6 Avagadro's Law and Gas Volumes

Empirical and Molecular Formula

Specification Reference S1.4.4

Quick Notes

  • Empirical formula = simplest whole number ratio of atoms in a compound.
  • Molecular formula = actual number of atoms of each element in a molecule.
  • The law of definite proportions states that a pure compound always has the same ratio of elements by mass.
  • To determine an empirical formula from mass data:
    • Convert mass (or %) of each element to moles.
    • Divide each mole value by the smallest to get a ratio.
    • Multiply to get whole numbers if necessary.
  • Key definitions:
    • Empirical formula: lowest whole number ratio.
    • Molecular formula: actual number of atoms (may be a multiple of empirical formula).

Full Notes

A chemical formula tells us the types and number of atoms in a substance. Some substances are made up of individual molecules (like H2O or C6H12O6), while others consist of ions or atoms in fixed proportions, like NaCl or Fe2O3.

Empirical vs Molecular Formula

The empirical formula gives the simplest whole number ratio of elements in a compound.

The molecular formula shows the actual number of atoms of each element in one molecule of the compound.

ExampleHydrogen peroxide has the molecular formula H2O2, but its empirical formula is HO. Glucose has the molecular formula C6H12O6, and its empirical formula is CH2O.

Determining the Empirical Formula from Mass Data

To find the empirical formula from mass or percentage composition, follow these steps:

Step 1: Convert each element’s mass (or percentage) to moles
If given %, treat the total as 100 g
Use: moles = mass ÷ molar mass
Divide mole amounts by the smallest value
Multiply up to get only whole numbers if needed (i.e. x everything by 2)

Worked Example

Find the empirical formula for a compound with a composition by mass of C 52.2%, H 13.0% and O 34.8%.

  1. Assume 100 g total
    C = 52.2 g → 52.2 ÷ 12.01 = 4.35 mol
    H = 13.0 g → 13.0 ÷ 1.008 = 12.9 mol
    O = 34.8 g → 34.8 ÷ 16.00 = 2.18 mol
  2. Divide by smallest mole value (2.18)
    C = 4.35 ÷ 2.18 = 2.00
    H = 12.9 ÷ 2.18 = 5.92 ≈ 6.00
    O = 2.18 ÷ 2.18 = 1.00
  3. Write empirical formula
    C2H6O
IB Chemistry worked example showing C 52.2%, H 13.0%, O 34.8% converted to moles and ratios to give empirical formula C2H6O.

Finding the Molecular Formula

If the molar mass of the compound is known, you can find the molecular formula:

Worked Example

The empirical formula of a compound is CH2O and its molar mass is 180 g/mol. Find the molecular formula.

  1. Calculate the empirical formula mass
    (1 × 12.0) + (2 × 1.0) + (1 × 16.0) = 30.0 g/mol
  2. Find n
    n = 180 ÷ 30.0 = 6
  3. Multiply empirical formula by n
    6 × CH2O = C6H12O6
IB Chemistry worked example showing CH2O empirical mass of 30.0, n = 180/30 = 6, giving molecular formula C6H12O6.

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Matt’s exam tip

Always check that your final empirical formula makes chemical sense. It should use whole numbers and reflect what’s possible based on the elements involved. If your ratio gives 1.5 or 1.33, multiply all ratios to clear the decimal.

Summary

Linked Questions

Tool 1 – Linked Course Question

How can experimental data on mass changes in combustion reactions be used to derive empirical formulas?

In combustion analysis, a compound containing carbon and hydrogen (and sometimes oxygen) is burned in excess oxygen, and the masses of CO₂ and H₂O produced are measured. From these data, the moles of C and H in the original sample can be determined. If oxygen is present in the compound, its mass is found by subtraction from the total. The mole ratios of the elements are then simplified to the smallest whole numbers, giving the empirical formula – the simplest ratio of atoms in the compound.


Nature of Science, Tool 3, Structure 3.2 – Linked Course Question

What is the importance of approximation in the determination of an empirical formula?

Experimental data rarely produce perfect whole-number ratios of atoms because of measurement uncertainty and rounding in molar mass calculations. Approximations are used to interpret these near-ratios – for example, a mole ratio of 1.99:1 is taken as 2:1. Recognising when to round and when to multiply to achieve whole-number ratios is an essential skill that balances mathematical precision with chemical reasoning. This highlights the role of approximation as a necessary and accepted part of scientific data interpretation.