Atomic Structure and Isotopes: Relative Mass and Mass Spectrometry

This guide covers 2.1.1, Atomic structure and isotopes, from Module 2, Foundations in chemistry, of OCR A Level Chemistry A H432 – the subtopic that comes immediately before Compounds, formulae and equations (2.1.2) and Amount of substance (2.1.3) in the same module. It establishes the atomic-level vocabulary that the rest of the course, including mole calculations, relies on.

Isotopes

You need to be able to explain isotopes as atoms of the same element with different numbers of neutrons and different masses. Because isotopes of an element have identical proton and electron numbers, they behave almost identically in chemical reactions – only their physical properties, particularly mass and anything mass-dependent, differ meaningfully.

Atomic structure from atomic number, mass number and charge

You need to be able to state atomic structure in terms of the numbers of protons, neutrons and electrons for atoms and ions, given the atomic number, mass number and any ionic charge. Working this out systematically: protons equal the atomic number; neutrons equal the mass number minus the atomic number; electrons equal the atomic number, adjusted by subtracting the charge for a cation or adding it for an anion. This subtopic is also explicitly linked to two “How Science Works” strands in the specification: that different models of atomic structure explain different phenomena (for example, the Bohr model explaining periodic properties), and that accepted models of atomic structure have changed over time as evidence has been used to accept or reject particular models.

Relative isotopic mass and relative atomic mass

You need to be able to give the definitions of two closely related but distinct terms, both based on 1/12th of the mass of a carbon-12 atom as the standard for atomic masses:

• Relative isotopic mass – the mass of a single isotope, compared with 1/12th the mass of carbon-12.
• Relative atomic mass – the weighted mean mass of all of an element’s naturally occurring isotopes, compared with 1/12th the mass of carbon-12.

The specification requires these definitions to be given explicitly, so learn the exact wording – “weighted mean” is the detail most often dropped when candidates describe relative atomic mass from memory.

Mass spectrometry

You need to be able to use mass spectrometry in determining the relative isotopic masses and relative abundances of an element’s isotopes, and in calculating relative atomic mass from those relative abundances. The specification is explicit that knowledge of how a mass spectrometer physically works is not required at this stage – what is required is being able to use mass spectrum data (masses and relative abundances of each isotope peak) to calculate a weighted mean relative atomic mass. This calculation is also limited to ions with single charges only.

The calculation itself is a weighted average: multiply each isotope’s mass by its percentage abundance, sum these products, then divide by 100. For example, for an element with 75% of one isotope at mass 35 and 25% of another at mass 37, the relative atomic mass is (75 × 35 + 25 × 37) ÷ 100 = 35.5.

Relative molecular mass and relative formula mass

You need to be able to use the terms relative molecular mass (Mr) and relative formula mass, and calculate both from relative atomic masses. The specification draws a precise distinction here: relative molecular mass is the term used for simple molecules, while relative formula mass is used for compounds with giant structures (such as ionic lattices, which do not exist as discrete molecules). Formal definitions of these two terms are not required, but knowing which term applies to which type of substance is expected, and questions may specifically test whether you have applied the correct term to a giant ionic structure rather than defaulting to “molecular mass” out of habit.

A worked example

Chlorine has two naturally occurring isotopes: chlorine-35, relative abundance 75.8%, and chlorine-37, relative abundance 24.2%. To find the relative atomic mass, calculate the weighted mean: (35 × 75.8 + 37 × 24.2) ÷ 100 = (2653 + 895.4) ÷ 100 = 35.5 (to 3 significant figures). This matches the value on the periodic table, and shows why chlorine’s relative atomic mass is not a whole number even though each individual isotope has a whole-number mass number – the weighting by abundance is what produces the familiar 35.5 figure. A related ion calculation: for the ³⁵Cl⁻ ion, the mass number 35 and atomic number 17 give 17 protons, 35 − 17 = 18 neutrons, and 17 + 1 = 18 electrons (one extra electron for the single negative charge).

How to approach it

Practise the proton/neutron/electron calculation from atomic number, mass number and ionic charge until it is instant, since it underlies every subsequent stoichiometry calculation in Module 2. Memorise the precise wording of the relative isotopic mass and relative atomic mass definitions – both reference 1/12th the mass of carbon-12, and the “weighted mean” phrase in the relative atomic mass definition is specifically what distinguishes it from relative isotopic mass. For mass spectrometry questions, focus on the weighted-average calculation using abundance data rather than trying to recall the instrument’s mechanism, since the specification explicitly excludes that mechanism from what you need to know. Finally, get in the habit of checking whether a substance is molecular or giant-structured before deciding whether to call its mass “relative molecular mass” or “relative formula mass” – ionic compounds such as sodium chloride should always use the latter term.

Official syllabus

OCR, Cambridge OCR Level 3 Advanced GCE in Chemistry A (H432) specification, Version 3.1, May 2026: official specification PDF, section 2.1.1 “Atomic structure and isotopes”. Verified 2026-09-02.