Size and mass of atoms — AQA GCSE Combined Science
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Size and mass of atoms explained
Atoms are the basic building blocks of matter, but they are far too small to see directly.
Read the full explanation
Their radius is about 0.1 nm, which is 1 x 10⁻¹⁰ m. To appreciate this scale, imagine lining up atoms across a 1 mm gap: about 5 million atoms would fit. A nanometre is one billionth of a metre, so 0.1 nm is one ten-billionth of a metre. This tiny size explains why atoms were not observed until powerful microscopes were developed. When comparing atomic sizes, remember that the value is a typical radius, not a fixed diameter, and different elements have slightly different atomic radii. In calculations, convert nanometres to metres by multiplying by 1 x 10⁻⁹, so 0.1 nm becomes 1 x 10⁻¹⁰ m.
The radius of a nucleus is less than 1/10 000 of that of the atom (about 1 x 10⁻¹⁴ m).
The nucleus sits at the centre of an atom and is much smaller than the atom itself. Its radius is less than 1/10 000 of the atomic radius, roughly 1 x 10⁻¹⁴ m. If an atom were scaled up to the size of a sports stadium, the nucleus would be about the size of a pea in the middle. This means most of an atom is empty space, with electrons occupying the outer regions. To compare, divide the atomic radius by the nuclear radius: 1 x 10⁻¹⁰ m divided by 1 x 10⁻¹⁴ m gives 10 000, so the atom is about 10 000 times wider than its nucleus. The nucleus contains protons and neutrons, and although it is tiny, it contains nearly all the mass of the atom.
Almost all of the mass of an atom is in the nucleus.
An atom is mostly empty space. The nucleus, a tiny central region containing protons and neutrons, carries almost all of the atom's mass because each proton and neutron has a relative mass of 1, while each electron has a relative mass of about 1/1836. The electrons occupy energy levels around the nucleus and contribute very little mass. For example, in a carbon-12 atom the nucleus contains 6 protons and 6 neutrons, giving a relative mass of 12, whereas its 6 electrons add only about 0.003. This explains why the mass number equals the number of protons plus neutrons and why relative atomic mass is dominated by nuclear particles.
The relative masses of protons, neutrons and electrons are:
Relative mass compares the mass of a particle with a standard. In GCSE Combined Science, a proton has a relative mass of 1, a neutron has a relative mass of 1, and an electron has a relative mass of about 1/1836, often treated as negligible. These values explain why the mass number of an atom is the sum of its protons and neutrons, and why electrons barely affect atomic mass. For example, a sodium atom with 11 protons and 12 neutrons has a mass number of 23; its 11 electrons add only about 0.006 to the relative mass. The values are relative, not actual masses in kilograms.
The sum of the protons and neutrons in an atom is its mass number.
The mass number counts the particles packed into an atom's nucleus: protons and neutrons together. Electrons are ignored because their mass is negligible compared with nucleons. For sodium, 11 protons plus 12 neutrons gives a mass number of 23, written as ²³Na or shown top-left of the symbol. To find neutrons, subtract the atomic number from the mass number: 23 − 11 = 12. The atomic number identifies the element and equals the proton count; the mass number identifies the particular atom. Because the mass number is a count of nucleons, it is a whole number with no unit. Isotopes of an element share the same atomic number but differ in mass number.
Atoms of the same element can have different numbers of neutrons; these atoms are called isotopes of that element.
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, so they have the same atomic number but different mass numbers. Chlorine, for example, has isotopes with mass numbers 35 and 37 (³⁵Cl and ³⁷Cl): both have 17 protons, but one has 18 neutrons and the other 20. Because proton number fixes the element, isotopes share chemical properties; differing neutron numbers change mass and some physical behaviour. Relative atomic mass is the weighted mean of isotope masses, so it need not be a whole number. Isotopes may be stable or radioactive.
Atoms can be represented as shown in this example:
This statement introduces the standard symbolic representation of an atom, often called nuclide notation. In this convention the mass number is written as a superscript before the element symbol and the atomic number as a subscript before it, for example ²³Na₁₁ or, more commonly in AQA materials, ²³₁₁Na. The mass number counts protons plus neutrons in the nucleus, while the atomic number counts protons and defines the element. For a neutral atom, the number of electrons equals the number of protons, so the symbol also tells you the electron count. For example, ²³₁₁Na has 11 protons, 12 neutrons and 11 electrons. The representation is a compact way of showing the composition of any atom or ion, and it underpins calculations of subatomic particle numbers.
Students should be able to calculate the numbers of protons, neutrons and electrons in an atom or ion, given its atomic number and mass number.
This statement requires you to use two numbers to find the three subatomic particle counts. The atomic number equals the number of protons and identifies the element. The mass number equals protons plus neutrons, so neutrons are found by subtracting the atomic number from the mass number. In a neutral atom, electrons equal protons. In an ion, the charge tells you the electron change: a positive charge means electrons have been lost, so electrons equal protons minus the charge; a negative charge means electrons have been gained, so electrons equal protons plus the charge. For example, ²⁴₁₂Mg²⁺ has 12 protons, 12 neutrons and 10 electrons, while ³²₁₆S²⁻ has 16 protons, 16 neutrons and 18 electrons.
Students should be able to relate size and scale of atoms to objects in the physical world.
Atoms are extremely small, with a typical radius of about 1 × 10⁻¹⁰ m. To relate this to the physical world, consider a full stop about 1 mm across: roughly 5 million atoms could sit side by side across it. Another comparison is a human hair, which is roughly 1 million atoms wide. These comparisons show that atoms are far smaller than everyday objects. The nucleus is smaller still, with a radius of about 1 × 10⁻¹⁴ m, which is less than 1/10,000th of the atomic radius, meaning most of an atom is empty space. Scale diagrams and powers of ten help keep these numbers manageable.
Your focus
- State the approximate radius of an atom as 0.1 nm.
- Convert 0.1 nm into metres using standard form.
- Describe how the small size of atoms compares with everyday objects.
Show all 27 objectives
- State that the nuclear radius is less than 1/10 000 of the atomic radius.
- Give the approximate nuclear radius as 1 x 10⁻¹⁴ m.
- Explain that most of an atom is empty space because the nucleus is so small.
- Describe the location and relative mass of protons, neutrons and electrons in an atom.
- Explain why almost all of the mass of an atom is concentrated in the nucleus.
- Use relative masses to compare the contribution of nuclear particles and electrons to atomic mass.
- Recall the relative masses of protons, neutrons and electrons.
- Explain why electrons are treated as having negligible mass in atomic mass calculations.
- Apply relative masses to calculate the mass number of an atom from its proton and neutron numbers.
- State that mass number is the total number of protons and neutrons in an atom.
- Calculate the number of neutrons from mass number and atomic number.
- Interpret nuclide notation to identify mass number and atomic number.
- Define isotopes as atoms of the same element with different numbers of neutrons.
- Explain why isotopes of an element have the same chemical properties.
- Relate isotopic abundance to the weighted mean relative atomic mass of an element.
- Label the mass number and atomic number positions in a given atomic symbol.
- Calculate the number of neutrons from a supplied mass number and atomic number.
- Interpret a supplied atomic symbol to state the number of protons, neutrons and electrons.
- Calculate neutron number from mass number and atomic number.
- Determine electron number for a neutral atom and for a charged ion.
- Explain how the sign and magnitude of an ionic charge change the electron count.
- State the approximate radius of a typical atom in metres.
- Compare the size of an atom with a named everyday object using a calculation or scale diagram.
- Explain why the nucleus is described as very small compared with the whole atom.
Size and mass of atoms exam tips
Marking Points
- States that the radius of an atom is about 0.1 nm.
- Converts 0.1 nm to 1 x 10⁻¹⁰ m correctly.
- Recognises that atoms are extremely small and cannot be seen with the naked eye.
- Uses the scale of 1 nm = 1 x 10⁻⁹ m to convert between units.
- Describes the approximate size of an atom relative to everyday objects, such as about 5 million atoms across 1 mm.
- States that the radius of a nucleus is less than 1/10 000 of the radius of the atom.
- Gives the approximate nuclear radius as 1 x 10⁻¹⁴ m.
- Calculates the ratio of atomic radius to nuclear radius as about 10 000.
- Explains that most of the atom is empty space because the nucleus is so small.
- Recognises that the nucleus contains most of the atom's mass despite its small size.
- State that the nucleus contains protons and neutrons, which are the particles that give an atom almost all of its mass.
- Compare relative masses: a proton and a neutron each have a relative mass of 1, whereas an electron has a relative mass of about 1/1836.
- Explain that electrons occupy energy levels around the nucleus and contribute a negligible fraction of the total atomic mass.
- Use the carbon-12 example: 6 protons plus 6 neutrons give a relative mass of 12, while 6 electrons add only about 0.003.
- Link the idea to mass number: mass number equals the number of protons plus the number of neutrons, so it reflects nuclear mass.
- State that a proton has a relative mass of 1.
- State that a neutron has a relative mass of 1.
- State that an electron has a relative mass of about 1/1836, which is often treated as negligible.
- Explain that these relative masses are comparisons, not actual masses in kilograms.
- Use the values to calculate mass number as the sum of protons and neutrons, ignoring the electron contribution.
- Mass number = number of protons + number of neutrons in the nucleus.
- Electrons are not included because their mass is negligible relative to protons and neutrons.
- Atomic number = number of protons; for a neutral atom it also equals the number of electrons.
- Number of neutrons = mass number − atomic number.
- Mass number is a whole-number count of nucleons and carries no unit.
- In nuclide notation, the mass number is written as a superscript before the symbol, for example ²³Na.
- Isotopes have the same number of protons, so they are atoms of the same element.
- Isotopes have different numbers of neutrons and therefore different mass numbers.
- The atomic number is the same for all isotopes of an element.
- Chemical properties are essentially the same because electron arrangement is the same.
- Relative atomic mass is a weighted mean of the masses of the isotopes, so it may not be a whole number.
- Isotopes of an element may include stable and radioactive forms.
- Recognises that the superscript before the symbol is the mass number, equal to protons plus neutrons.
- Recognises that the subscript before the symbol is the atomic number, equal to the number of protons.
- States that in a neutral atom the number of electrons equals the number of protons.
- Uses a given symbol such as ²³₁₁Na to deduce 11 protons, 12 neutrons and 11 electrons.
- Explains that the atomic number determines the element and that the mass number distinguishes isotopes of that element.
- States that the atomic number equals the number of protons.
- Calculates neutrons by subtracting the atomic number from the mass number.
- States that electrons equal protons in a neutral atom.
- Adjusts the electron count for an ion by subtracting the positive charge or adding the magnitude of the negative charge.
- Applies the method correctly to a supplied atom or ion, for example ²⁴₁₂Mg²⁺ giving 12 protons, 12 neutrons and 10 electrons.
- States that a typical atom has a radius of about 1 × 10⁻¹⁰ m.
- Uses a familiar object, such as a full stop or a human hair, to compare with atomic size.
- Explains that a human hair is roughly 1 million atoms wide.
- Describes the nucleus as having a radius of about 1 × 10⁻¹⁴ m.
- Recognises that most of an atom is empty space between the nucleus and the electrons.
Examiner Tips
- 💡Learn the conversion 1 nm = 1 x 10⁻⁹ m and practise converting 0.1 nm into metres.
- 💡When asked to compare sizes, quote the atomic radius as about 0.1 nm or 1 x 10⁻¹⁰ m.
- 💡Use standard form correctly in calculations and show your working clearly.
- 💡Memorise the approximate values: atomic radius about 1 x 10⁻¹⁰ m and nuclear radius about 1 x 10⁻¹⁴ m.
- 💡When comparing sizes, divide the atomic radius by the nuclear radius to show the atom is about 10 000 times wider.
- 💡Use the idea of empty space to explain why atoms are mostly empty.
- 💡Use the phrase 'almost all of the mass' rather than 'all of the mass', because electrons do contribute a very small amount.
- 💡When explaining, quote relative masses: proton 1, neutron 1, electron about 1/1836.
- 💡Link your answer to the structure of the atom: a small, dense nucleus surrounded by electrons in energy levels.
- 💡Learn the values: proton 1, neutron 1, electron about 1/1836.
- 💡When calculating mass number, add protons and neutrons only; do not include electrons.
- 💡If asked to compare, state that protons and neutrons have similar relative masses, while electrons are much lighter.
- 💡Write the equation mass number = protons + neutrons before substituting values, so the arithmetic is clear.
- 💡Check that your neutron count is a whole number; a fractional answer signals a slip.
- 💡When reading nuclide notation, identify the superscript as mass number and the subscript as atomic number before calculating.
- 💡Define isotopes using both 'same protons' and 'different neutrons' to secure the mark.
- 💡Use a named example such as chlorine-35 and chlorine-37 to make the explanation concrete.
- 💡If asked to compare, state clearly what is the same and what is different rather than describing each isotope separately.
- 💡Always write the mass number above and the atomic number below, aligned to the left of the element symbol.
- 💡When asked for particle numbers, show the subtraction mass number minus atomic number for neutrons.
- 💡Check whether the species is neutral or an ion before stating the electron count.
- 💡Write the three quantities in a consistent order, for example protons, neutrons, electrons, to avoid mixing them up.
- 💡Show the neutron calculation explicitly as mass number minus atomic number.
- 💡For ions, state whether electrons are gained or lost before giving the final number.
- 💡Quote the approximate atomic radius as 1 × 10⁻¹⁰ m and link it to a named everyday object.
- 💡When comparing nucleus and atom, give both approximate radii so the scale difference is clear.
Common Mistakes
- Writing 0.1 nm as 1 x 10⁻⁹ m instead of 1 x 10⁻¹⁰ m. Correction: 0.1 nm = 0.1 x 10⁻⁹ m = 1 x 10⁻¹⁰ m.
- Confusing radius with diameter. Correction: the stated value is a radius; the diameter would be about 0.2 nm.
- Using the word 'atom' to mean the nucleus. Correction: the nucleus is much smaller and is found at the centre of the atom.
- Writing the nuclear radius as 1 x 10⁻¹⁰ m instead of 1 x 10⁻¹⁴ m. Correction: the atomic radius is about 1 x 10⁻¹⁰ m and the nuclear radius is about 1 x 10⁻¹⁴ m.
- Thinking the nucleus occupies most of the atom's volume. Correction: the nucleus is tiny; most of the atom is empty space.
- Confusing the ratio 1/10 000 with 1/1000. Correction: the nuclear radius is less than one ten-thousandth of the atomic radius.
- Thinking that electrons make a large contribution to atomic mass because they move quickly; correct this by comparing relative masses, where an electron is about 1/1836 of a proton.
- Believing that the nucleus occupies most of the atom's volume; correct this by stating that the nucleus is tiny compared with the whole atom, even though it contains almost all the mass.
- Confusing mass number with relative atomic mass; correct this by noting that mass number counts protons and neutrons, while relative atomic mass is a weighted mean that also accounts for isotopes.
- Writing that an electron has a relative mass of 0; correct this by saying it is about 1/1836, which is very small but not zero.
- Giving the relative mass of a neutron as 0; correct this by stating that a neutron has a relative mass of 1, like a proton.
- Confusing relative mass with actual mass in kilograms; correct this by explaining that relative mass is a ratio compared with a standard, so it has no units.
- Including electrons in the mass number: correct by stating that only protons and neutrons are counted because electron mass is negligible.
- Confusing mass number with relative atomic mass: correct by noting that mass number is a whole-number nucleon count, whereas relative atomic mass is a weighted mean that may not be whole.
- Subtracting incorrectly to find neutrons, for example doing atomic number − mass number: correct by always using neutrons = mass number − atomic number.
- Saying isotopes have different numbers of protons: correct by stating that proton number is identical and only neutron number differs.
- Claiming isotopes have completely different chemical properties: correct by explaining that identical electron arrangements give essentially the same chemistry.
- Treating relative atomic mass as a simple average of mass numbers: correct by using the weighted mean that accounts for isotopic abundance.
- Writing the mass number as a subscript and the atomic number as a superscript; correct this by remembering that the larger number is the mass number and sits above the atomic number.
- Assuming the number of neutrons equals the atomic number; correct this by calculating neutrons as mass number minus atomic number.
- Treating the mass number as the number of electrons; correct this by using the atomic number for protons and electrons in a neutral atom.
- Adding the charge to the electron count for a positive ion; correct this by subtracting the charge because cations have lost electrons.
- Using the mass number as the proton count; correct this by using the atomic number for protons.
- Forgetting to adjust electrons for an ion and giving the neutral-atom value; correct this by checking the sign and magnitude of the charge before finalising the electron number.
- Saying an atom has a diameter of 1 × 10⁻¹⁰ m; correct this by stating it is the radius.
- Confusing the radius of the nucleus with the atom; correct this by noting the nucleus radius is about 1 × 10⁻¹⁴ m.
- Writing 10⁻¹⁰ as 10-10; correct this by using the raised exponent 10⁻¹⁰.