Relative Atomic MassOxford AQA IGCSE Chemistry: Revision notes

Section 1

Why do we need relative masses?

Atoms are far too small and light to weigh individually in grams. Instead, scientists compare the masses of atoms relative to one another, using the carbon-12 isotope as the standard reference point.

This comparison gives the relative atomic mass ( $A_r$ ) of each element.

Key termsrelative atomic mass

Section 2

Defining relative atomic mass

Relative atomic mass ( $A_r$ ) is defined as an average value for the isotopes of an element, compared with the ¹²C isotope.

Carbon-12 is defined as having a mass of exactly 12
All other atoms' masses are compared against this standard
Because most elements exist as a mixture of isotopes in fixed natural proportions, $A_r$ is a weighted average, not the mass of any single atom

Key terms¹²Cweighted average

Common mistake

Do not assume relative atomic mass is always a whole number — because it is an average across isotopes, it is often a decimal, e.g. chlorine is 35.5.

Section 3

Relative masses of subatomic particles

Relative atomic mass depends on the masses of the particles that make up the nucleus:

Particle	Relative mass
Proton	1
Neutron	1
Electron	Very small (negligible)

Because electrons have negligible mass compared with protons and neutrons, the mass of an atom is essentially determined by the number of protons and neutrons in its nucleus.

Section 4

What you are not expected to do

At this level, you are not expected to calculate relative atomic masses from isotopic abundance data (e.g. combining percentage abundances and isotope masses into a weighted-average calculation). You only need to know what relative atomic mass represents and that it is an average value.

Exam tip

If a question gives isotope abundance data, focus on explaining the concept (why Ar is an average) rather than attempting a calculation, unless specifically asked to calculate.

Must Know

Relative atomic mass ( $A_r$ ) compares an element's atoms with the ¹²C isotope
$A_r$ is an average value across all naturally occurring isotopes of an element
Protons and neutrons each have a relative mass of 1; electrons have negligible mass
$A_r$ is often not a whole number because it is a weighted average (e.g. chlorine = 35.5)
You are not expected to calculate $A_r$ from isotopic abundance data

That's the notes covered.

Carry on to the next subtopic.