Understanding Relative Atomic Mass (Ar) in OCR A-Level Chemistry
- Oct 13, 2025
- 5 min read
Updated: Jul 25
By Paul Morgan
Relative atomic mass (Ar) is one of the first calculations you'll encounter in OCR A-Level Chemistry.
Although the concept seems straightforward, many students confuse relative atomic mass, mass number and isotopes, leading to avoidable mistakes in exams.
Understanding relative atomic mass is essential because it forms the foundation for many other topics, including moles, empirical formulae, reacting masses, limiting reagents and percentage yield calculations.
In this guide you'll learn:
What relative atomic mass actually means
Why isotopes affect relative atomic mass
How to calculate relative atomic mass using isotope abundance
Common OCR exam mistakes
Worked examples and practice questions
By the end of this guide, you'll be able to calculate relative atomic mass confidently and understand why it is used throughout the OCR A-Level Chemistry specification.
What Is Relative Atomic Mass?
Relative atomic mass (Ar) is the weighted average mass of an atom of an element compared with one-twelfth of the mass of a carbon-12 atom.
Carbon-12 has been chosen as the international standard and is assigned a relative atomic mass of exactly 12.
Every other element is compared with this standard.
Because atoms are incredibly small, chemists use relative masses rather than measuring their actual masses.
For example, oxygen has a relative atomic mass of approximately 16.
This means an average oxygen atom has a mass approximately sixteen times greater than one-twelfth of the mass of a carbon-12 atom.
Why Is Relative Atomic Mass Important?
Relative atomic mass is used throughout OCR A-Level Chemistry.
You'll use it when calculating:
Moles
Reacting masses
Empirical formulae
Molecular formulae
Percentage yield
Atom economy
Without a good understanding of relative atomic mass, many quantitative chemistry calculations become much more difficult.
Why Don't All Elements Have Whole Number Relative Atomic Masses?
Many students notice that chlorine has a relative atomic mass of 35.5 and wonder why it isn't simply 35 or 37.
The answer is isotopes.
Isotopes are atoms of the same element that contain the same number of protons but different numbers of neutrons.
This means isotopes have different masses.
Most elements occur naturally as mixtures of isotopes.
The published relative atomic mass is therefore a weighted average, taking into account both:
the mass of each isotope
its natural abundance
How to Calculate Relative Atomic Mass
Calculating relative atomic mass is straightforward once you understand the method.
Step 1
Write down the mass of each isotope.
2
Convert the percentage abundances into decimals.
For example:
75% = 0.75
25% = 0.25
Step 3
Multiply each isotope mass by its abundance.
Step 4
Add the results together.
The answer is the relative atomic mass.
Worked Example – Chlorine
Chlorine occurs naturally as two main isotopes.
Chlorine-35 (75%)
Chlorine-37 (25%)
Calculate the relative atomic mass.
Convert the percentages into decimals.
75% = 0.75
25% = 0.25
Now calculate the weighted average.
(35 × 0.75) + (37 × 0.25)
= 26.25 + 9.25
= 35.5
Therefore the relative atomic mass of chlorine is:
Ar = 35.5
Worked Example – Copper
Copper also exists naturally as two common isotopes.
Copper-63 (69%)
Copper-65 (31%)
Calculate the relative atomic mass.
Convert the percentages.
69% = 0.69
31% = 0.31
Now calculate the weighted average.
(63 × 0.69) + (65 × 0.31)
= 43.47 + 20.15
= 63.62
Therefore:
Ar(Cu) = 63.6
Common OCR Exam Mistakes
Even students who understand isotopes often lose marks through simple mistakes.
Mistake 1
Confusing mass number with relative atomic mass.
Mass number refers to one isotope.
Relative atomic mass is the weighted average of all naturally occurring isotopes.
Mistake 2
Simply averaging the isotope masses.
For chlorine:
(35 + 37) ÷ 2 = 36
This is incorrect because chlorine-35 is much more abundant than chlorine-37.
Always calculate a weighted average.
Mistake 3
Forgetting to convert percentages into decimals.
75%
must become
0.75
before carrying out the calculation.
Mistake 4
Using relative atomic mass instead of relative formula mass.
Remember:
Relative atomic mass is used for elements.
Relative formula mass is used for compounds.
OCR Exam Tip
Whenever an OCR question gives isotope abundances, don't rush into the calculation.
First ask yourself:
Is this asking for a simple average or a weighted average?
If isotope abundances are given, the answer is always a weighted average.
Ignoring the abundances is one of the most common mistakes students make.
Practice Question
Bromine occurs naturally as two isotopes.
Bromine-79 (51%)
Bromine-81 (49%)
Calculate the relative atomic mass of bromine.
Try solving the question before looking at the answer.
Answer
Convert the percentages into decimals.
51% = 0.51
49% = 0.49
Now calculate the weighted average.
(79 × 0.51) + (81 × 0.49)
= 40.29 + 39.69
= 79.98
Therefore:
Relative atomic mass of bromine = 80.0
How Relative Atomic Mass Is Used Later in OCR Chemistry
Relative atomic mass isn't just an isolated topic.
It appears throughout the OCR specification.
You'll use it when calculating:
Relative formula mass (Mr)
Moles
Empirical formulae
Molecular formulae
Percentage yield
Limiting reagents
Atom economy
Mastering relative atomic mass now will make these later calculations much easier.
Frequently Asked Questions
Why isn't chlorine's relative atomic mass exactly 35 or 37?
Because natural chlorine contains a mixture of chlorine-35 and chlorine-37 atoms.
The published value is the weighted average of both isotopes.
Why isn't relative atomic mass always a whole number?
Because most elements occur naturally as mixtures of isotopes with different abundances.
The weighted average usually produces a decimal value.
What is the difference between mass number and relative atomic mass?
Mass number refers to one isotope and is always a whole number.
Relative atomic mass is the weighted average of all naturally occurring isotopes.
Why is carbon-12 used as the standard?
Carbon-12 has been chosen internationally as the reference isotope.
Its mass is defined as exactly 12, allowing every other element to be compared with it.
Key Takeaways
Before your OCR Chemistry exam, remember these five rules:
Relative atomic mass compares atoms with carbon-12.
It is the weighted average of naturally occurring isotopes.
Isotope abundance must always be included in calculations.
Relative atomic mass is used throughout mole and stoichiometry calculations.
Never confuse relative atomic mass with mass number.
Master this topic and you'll find later calculations involving moles, reacting masses and limiting reagents much easier.
Free OCR A-Level Chemistry Guides
If you're studying OCR A-Level Chemistry and want to avoid the most common mistakes that hold students back, download my free guides:
Year 12 Students:
4 Mistakes That Cause Strong GCSE Students To Struggle In Year 12 Chemistry
Year 13 Students:
4 Mistakes Keeping Capable OCR Chemistry Students Stuck At Grade B Or Below
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