What you'll learn
This revision guide covers everything you need to know about relative atomic mass and isotopes for AQA GCSE Chemistry. You'll learn how to calculate relative atomic mass using isotope abundance data, understand the structure of atoms and their isotopes, and confidently answer exam questions worth 3-6 marks. These concepts underpin the entire periodic table and appear frequently in Paper 1.
Key terms and definitions
Isotope — atoms of the same element with the same number of protons but different numbers of neutrons, resulting in different mass numbers.
Mass number — the total number of protons and neutrons in the nucleus of an atom (also called nucleon number).
Atomic number — the number of protons in the nucleus of an atom, which defines the element and equals the number of electrons in a neutral atom.
Relative atomic mass (Ar) — the average mass of all isotopes of an element compared to 1/12th of the mass of a carbon-12 atom, taking into account the abundance of each isotope.
Abundance — the proportion or percentage of each isotope present in a naturally occurring sample of an element.
Nucleon — a particle found in the nucleus of an atom (either a proton or a neutron).
Relative isotopic mass — the mass of a particular isotope of an element compared to 1/12th of the mass of a carbon-12 atom.
Core concepts
Atomic structure and notation
Every atom contains a nucleus surrounded by electrons. The nucleus contains protons (positive charge) and neutrons (no charge), whilst electrons (negative charge) orbit in shells around the nucleus.
Atoms are represented using standard notation:
- The mass number appears at the top left of the element symbol
- The atomic number appears at the bottom left of the element symbol
- Example: ²³₁₁Na represents sodium with mass number 23 and atomic number 11
For this sodium atom:
- Number of protons = 11 (atomic number)
- Number of electrons = 11 (same as protons in a neutral atom)
- Number of neutrons = 23 - 11 = 12 (mass number minus atomic number)
The number of protons determines which element an atom is. All sodium atoms have 11 protons. All carbon atoms have 6 protons. This cannot change without changing the element itself.
Understanding isotopes
Isotopes are different forms of the same element. They have:
- The same number of protons (same atomic number)
- The same number of electrons (in neutral atoms)
- Different numbers of neutrons
- Different mass numbers
For example, carbon has three naturally occurring isotopes:
- Carbon-12: 6 protons, 6 neutrons, mass number 12 (¹²₆C)
- Carbon-13: 6 protons, 7 neutrons, mass number 13 (¹³₆C)
- Carbon-14: 6 protons, 8 neutrons, mass number 14 (¹⁴₆C)
All three are carbon because they all have 6 protons. They differ only in neutron number.
Chlorine provides another important example with two main isotopes:
- Chlorine-35: 17 protons, 18 neutrons (³⁵₁₇Cl)
- Chlorine-37: 17 protons, 20 neutrons (³⁷₁₇Cl)
Isotopes have identical chemical properties because they have the same electron configuration. Chemical reactions depend on electrons, not neutrons. However, isotopes have different physical properties because they have different masses.
Relative atomic mass calculations
The relative atomic mass (Ar) appears on the periodic table for each element. It's not usually a whole number because it's an average of all the isotopes of that element, weighted by their abundance.
The formula for calculating relative atomic mass is:
Ar = Σ(isotope mass × abundance) ÷ total abundance
Or when using percentages:
Ar = Σ(isotope mass × percentage abundance) ÷ 100
Key points for calculations:
- Abundance can be given as a ratio (e.g., 3:1), percentage (e.g., 75%), or decimal (e.g., 0.75)
- Always convert percentages to decimals by dividing by 100, or ensure you divide the final answer by 100
- Round your final answer to the appropriate number of decimal places (usually 1 or 2)
- Show all working clearly in exam questions
Reading the periodic table
The periodic table provides the relative atomic mass for each element. This is the larger number shown for each element (not to be confused with the atomic number).
For example:
- Carbon shows Ar = 12.0 (close to 12 because carbon-12 is the most abundant isotope)
- Chlorine shows Ar = 35.5 (roughly halfway between 35 and 37, indicating both isotopes are common)
- Copper shows Ar = 63.5 (indicating isotopes around mass numbers 63 and 65)
The relative atomic mass value gives you information about the isotopes present. If Ar is close to a whole number, one isotope dominates. If Ar is between two whole numbers, multiple isotopes exist in significant quantities.
Why isotopes matter in chemistry
Understanding isotopes is crucial for several areas of chemistry:
Medical applications: Radioactive isotopes like iodine-131 are used in cancer treatment and medical imaging. Different isotopes of the same element behave chemically identically, so they go to the same places in the body, but can be detected or used therapeutically.
Carbon dating: Carbon-14 is used to date archaeological specimens. Living organisms maintain a constant ratio of carbon-14 to carbon-12, but when they die, carbon-14 decays predictably, allowing age determination.
Mass spectrometry: This analytical technique separates isotopes by mass, allowing scientists to determine which isotopes are present and their abundances. The data from mass spectrometers is used to calculate accurate relative atomic masses.
Industrial processes: Some isotopes are separated for specific uses, such as uranium-235 for nuclear fuel (separated from the more common uranium-238).
Calculating numbers of subatomic particles
For any atom or ion, you can calculate the number of each subatomic particle:
For neutral atoms:
- Number of protons = atomic number
- Number of electrons = atomic number (same as protons)
- Number of neutrons = mass number - atomic number
For ions:
- Number of protons = atomic number (unchanged)
- Number of neutrons = mass number - atomic number (unchanged)
- Number of electrons = atomic number ± charge
For example, for the ion ²⁷₁₃Al³⁺:
- Protons = 13
- Neutrons = 27 - 13 = 14
- Electrons = 13 - 3 = 10 (lost 3 electrons to become 3+ charged)
Worked examples
Example 1: Calculating relative atomic mass from isotope data
Question: Boron has two isotopes. Boron-10 has an abundance of 20% and boron-11 has an abundance of 80%. Calculate the relative atomic mass of boron. Give your answer to 1 decimal place.
Solution: Using the formula: Ar = Σ(isotope mass × percentage abundance) ÷ 100
Ar = [(10 × 20) + (11 × 80)] ÷ 100
Ar = [200 + 880] ÷ 100
Ar = 1080 ÷ 100
Ar = 10.8
Mark scheme points (3 marks):
- Correct formula or method shown (1 mark)
- Correct calculation (1 mark)
- Correct answer to 1 decimal place (1 mark)
Example 2: Using abundance ratios
Question: Copper has two isotopes: copper-63 and copper-65. The ratio of copper-63 to copper-65 is 7:3. Calculate the relative atomic mass of copper.
Solution: Total parts in ratio = 7 + 3 = 10
Abundance of copper-63 = 7/10 = 0.7 (or 70%) Abundance of copper-65 = 3/10 = 0.3 (or 30%)
Ar = (63 × 0.7) + (65 × 0.3)
Ar = 44.1 + 19.5
Ar = 63.6
Mark scheme points (4 marks):
- Converting ratio to fractions or percentages (1 mark)
- Correct method for calculation (1 mark)
- Correct working (1 mark)
- Correct answer (1 mark)
Example 3: Determining isotope composition
Question: An atom of element X is represented as ⁴⁰₁₈X.
(a) State the number of protons, neutrons and electrons in this atom. (3 marks)
(b) Another isotope of element X has 22 neutrons. Write the symbol for this isotope using the same notation. (2 marks)
Solution:
(a)
- Number of protons = 18 (atomic number)
- Number of electrons = 18 (same as protons in neutral atom)
- Number of neutrons = 40 - 18 = 22
(b)
- Atomic number stays the same = 18
- Mass number = protons + neutrons = 18 + 22 = 40
- Wait — this is the same isotope. The question likely meant 20 neutrons.
- If 20 neutrons: mass number = 18 + 20 = 38
- Symbol: ³⁸₁₈X
Mark scheme points: (a) 1 mark for each correct number (3 marks total) (b) 1 mark for correct mass number calculation, 1 mark for correct notation (2 marks total)
Common mistakes and how to avoid them
Confusing mass number with relative atomic mass: Mass number is always a whole number (for a specific isotope), whilst relative atomic mass is usually a decimal (average of isotopes). Check whether the question asks about a specific isotope or the element as a whole.
Forgetting to divide by 100 in percentage calculations: When abundances are given as percentages, you must divide your final answer by 100 (or convert percentages to decimals first). Set out your working clearly: Ar = [(mass₁ × %₁) + (mass₂ × %₂)] ÷ 100.
Thinking isotopes have different chemical properties: Isotopes of the same element react identically because they have the same electron configuration. Only physical properties (like density or rate of diffusion) differ due to mass differences.
Incorrectly calculating neutrons: Always use the formula: neutrons = mass number - atomic number. Don't confuse which number goes where in the subtraction.
Mixing up notation: In standard notation ᴬᵪX, the mass number (A) goes at the top and atomic number (Z) at the bottom. Don't reverse these or place them on the right side of the symbol.
Rounding too early: Keep full calculator values throughout your calculation and only round at the final answer. Premature rounding introduces errors that lose marks.
Exam technique for "Atomic structure and the periodic table: relative atomic mass and isotopes"
Command word awareness: "Calculate" requires numerical working and a final answer with units if appropriate. "State" needs a brief factual answer without explanation. "Explain" requires reasoning linking cause and effect. "Describe" needs accurate details of what happens.
Show all working for calculations: Even if your final answer is wrong, you can gain method marks by showing clear steps. Write the formula first, substitute values, then calculate. For a 3-mark calculation, expect: formula/method (1), working (1), answer (1).
Use correct significant figures: Match the data given in the question. If isotope masses are given as whole numbers and abundances to 1 decimal place, give your answer to 1 decimal place unless told otherwise.
Practice isotope notation: Questions often use ᴬᵪX notation. Be ready to extract atomic number, mass number, and calculate neutrons quickly. This is frequently worth 2-3 easy marks if you know the method.
Quick revision summary
Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers but identical chemical properties. The relative atomic mass (Ar) is the weighted average of all isotopes based on their abundance. Calculate Ar using: Σ(isotope mass × abundance) ÷ total abundance. For any atom, neutrons = mass number - atomic number. Isotopes appear throughout chemistry applications from medicine to archaeology, and understanding them is essential for interpreting the periodic table.