Mass number, atomic number and isotopes — AQA GCSE Combined Science
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Mass number, atomic number and isotopes explained
Every atom contains a nucleus of protons and neutrons, with electrons arranged in shells around it.
Read the full explanation
A proton carries a 1⁺ charge and an electron carries a 1⁻ charge, so when the numbers are equal the positive and negative charges cancel exactly. For example, carbon has six protons, so a neutral carbon atom also has six electrons, giving 6⁺ and 6⁻ and a net charge of zero. If the numbers differ, the particle is an ion, not an atom: losing an electron leaves a positive ion such as Na⁺, while gaining two electrons gives a negative ion such as O²⁻. The rule applies to atoms only, and the equal numbers explain why atoms are electrically neutral overall.
All atoms of a particular element have the same number of protons. The number of protons in an atom of an element is called its atomic number.
The atomic number defines the element. Every atom of a given element has the same proton count, and that count is the atomic number. For example, all carbon atoms have six protons, so carbon's atomic number is 6; all oxygen atoms have eight protons, so oxygen's atomic number is 8. Changing the proton number changes the element, so an atom with seven protons is nitrogen, not carbon. Atoms of the same element can differ in neutron number, forming isotopes, but the proton number stays fixed. In symbol notation the atomic number is written as the lower number, for example ₆C, and it also equals the number of electrons in a neutral atom.
The total number of protons and neutrons in an atom is called its mass number.
The mass number of an atom is the sum of its protons and neutrons, because these are the only nucleons with appreciable mass. Electrons contribute almost nothing to the mass of an atom, so they are excluded. For example, a sodium atom with 11 protons and 12 neutrons has a mass number of 11 + 12 = 23. The mass number is always a whole number and is written as a superscript before the symbol, as in ²³Na. It is not the same as relative atomic mass, which is a weighted mean of isotope masses and is not usually a whole number. To find the number of neutrons, subtract the atomic number from the mass number: neutrons = mass number − atomic number.
Atoms can be represented as shown in this example:
Atoms are represented using nuclide notation, which shows the mass number as a superscript and the atomic number as a subscript before the chemical symbol. For example, ²³₁₁Na represents a sodium atom with mass number 23 and atomic number 11. The atomic number tells you the number of protons, which identifies the element. The mass number tells you the total number of protons and neutrons. From these two numbers you can work out the number of neutrons: neutrons = mass number − atomic number. In ²³₁₁Na, the number of neutrons is 23 − 11 = 12. This notation is used for all atoms and is essential for identifying isotopes, which have the same atomic number but different mass numbers.
Atoms of the same element can have different numbers of neutrons; these atoms are called isotopes of that element.
Every atom of an element has the same number of protons, so the atomic number is fixed. The neutron number can vary, giving atoms of the same element with different mass numbers; these are isotopes. For example, chlorine has isotopes with mass numbers 35 and 37: each has 17 protons, but one has 18 neutrons and the other 20 neutrons. Because isotopes of an element have identical proton counts, they have identical electron counts in neutral atoms and therefore the same chemical properties; their physical properties and relative masses differ because the neutron number changes the mass. The relative atomic mass quoted for an element is a weighted mean of its isotopes' masses, so it is often not a whole number.
Atoms turn into positive ions if they lose one or more outer electron(s).
A neutral atom has equal numbers of protons and electrons. The outer electrons are the ones most easily removed because they are furthest from the nucleus and least strongly attracted. When an atom loses one or more of these outer electrons, the number of protons exceeds the number of electrons, so the overall charge becomes positive and the particle is a positive ion. For example, a sodium atom has 11 protons and 11 electrons; losing one outer electron leaves 11 protons and 10 electrons, giving a 1⁺ charge and the ion Na⁺. The nucleus is unchanged, so the proton number and the element identity stay the same; only the electron count and charge change.
Students should be able to relate differences between isotopes to differences in conventional representations of their identities, charges and masses.
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons. In conventional representations, the identity is the element symbol, the charge is shown as a superscript, and the mass is shown as a superscript mass number. For example, chlorine-35 and chlorine-37 are represented as ³⁵Cl and ³⁷Cl. Both have atomic number 17, so 17 protons and 17 electrons in a neutral atom, giving no overall charge. The mass number differs because the neutron number differs: 18 neutrons in ³⁵Cl and 20 neutrons in ³⁷Cl. If an isotope forms an ion, the charge is shown separately, such as ³⁵Cl⁻, which has 18 electrons and an overall charge of 1⁻. Thus differences in isotope identity, charge and mass are reflected in the mass number, charge superscript and element symbol.
Your focus
- State the relative charges of protons and electrons.
- Explain why an atom has no overall electrical charge.
- Identify whether a given particle is a neutral atom or an ion from its proton and electron numbers.
Show all 21 objectives
- Define atomic number as the number of protons in an atom.
- Explain why all atoms of an element have the same atomic number.
- Use atomic number to identify an element and distinguish it from its isotopes.
- State that mass number is the total number of protons and neutrons in an atom.
- Calculate the number of neutrons in an atom from its mass number and atomic number.
- Represent a nuclide using the correct notation with mass number as a superscript and atomic number as a subscript.
- Interpret nuclide notation to identify the number of protons and neutrons in an atom.
- Write nuclide notation correctly for a given atom or isotope.
- Explain how isotopes of the same element differ in mass number but not atomic number.
- Define isotopes as atoms of the same element with different numbers of neutrons.
- Determine the number of protons, neutrons and electrons in a given isotope from its notation.
- Explain why isotopes of an element have the same chemical properties but can differ in physical properties.
- Describe how a neutral atom becomes a positive ion by losing one or more outer electrons.
- Determine the charge of a positive ion from its proton and electron counts.
- Write the symbol of a positive ion with the charge shown as a superscript.
- Identify the number of protons, neutrons and electrons in a given isotope or ion from its conventional representation.
- Explain how differences in neutron number lead to different mass numbers for isotopes of the same element.
- Interpret and write conventional representations that show element identity, mass number and charge for atoms and ions.
Mass number, atomic number and isotopes exam tips
Marking Points
- State that a proton has a charge of 1⁺ and an electron has a charge of 1⁻, so equal numbers cancel.
- Explain that the nucleus contains protons and neutrons, and that neutrons have no charge.
- Apply the rule to a named element, for example carbon with six protons and six electrons.
- Distinguish a neutral atom from an ion, where electron number differs from proton number.
- Use the idea of overall charge as the sum of positive and negative charges being zero.
- Define atomic number as the number of protons in an atom of an element.
- State that all atoms of the same element have identical proton numbers.
- Use examples such as carbon with atomic number 6 and oxygen with atomic number 8.
- Explain that changing the proton number changes the element.
- Link atomic number to the number of electrons in a neutral atom.
- Mass number equals the number of protons plus the number of neutrons in an atom.
- Electrons are not included in the mass number because their mass is negligible compared with protons and neutrons.
- The mass number is a whole number and is written as a superscript before the chemical symbol, for example ²³Na.
- The number of neutrons can be calculated using: neutrons = mass number − atomic number.
- Mass number is not the same as relative atomic mass, which is a weighted mean and may not be a whole number.
- Nuclide notation places the mass number as a superscript and the atomic number as a subscript before the chemical symbol.
- The atomic number equals the number of protons and determines the element.
- The mass number equals the number of protons plus the number of neutrons.
- The number of neutrons is calculated as mass number minus atomic number.
- Isotopes of an element have the same atomic number but different mass numbers.
- State that isotopes are atoms of the same element, so they have the same number of protons and the same atomic number.
- State that isotopes differ in the number of neutrons, so they have different mass numbers.
- Calculate the neutron number of an isotope by subtracting the atomic number from the mass number.
- Explain that neutral isotopes of an element have the same electron arrangement, so their chemical properties are the same.
- Explain that different neutron numbers give different masses, so physical properties such as density or rate of diffusion can differ.
- Interpret or calculate a weighted mean relative atomic mass from given isotope abundances.
- State that a neutral atom has equal numbers of protons and electrons.
- State that losing electrons leaves more protons than electrons, producing an overall positive charge.
- Identify the electrons lost as outer electrons, and link this to their weaker attraction to the nucleus.
- Work out the charge from the numbers of protons and electrons, for example 11 protons and 10 electrons gives a 1⁺ charge.
- Write the ion symbol correctly, such as Na⁺ or Mg²⁺, showing the charge as a superscript.
- Recognise that the nucleus and proton number do not change when an atom forms a positive ion.
- Isotopes have the same number of protons but different numbers of neutrons, so their mass numbers differ.
- The conventional representation includes the element symbol, a mass number superscript and, for ions, a charge superscript.
- The mass number equals the number of protons plus the number of neutrons, so it distinguishes isotopes of the same element.
- The charge is determined by the difference between the number of protons and the number of electrons, not by the number of neutrons.
- For a neutral atom, the number of electrons equals the number of protons, so the overall charge is zero.
- For an ion, the charge superscript shows the overall charge, for example 1⁺, 2⁺, 1⁻ or 2⁻.
Examiner Tips
- 💡Quote the charges 1⁺ and 1⁻ when explaining why an atom is neutral.
- 💡Use a specific element with its proton and electron numbers to support your answer.
- 💡Check whether the question says atom or ion before applying the equal-numbers rule.
- 💡Write the definition of atomic number precisely, using the word protons.
- 💡Use a named element and its proton number to illustrate your answer.
- 💡If a symbol is given, read the lower number as the atomic number.
- 💡Read the question carefully to check whether it asks for mass number, atomic number or number of neutrons.
- 💡Show the addition or subtraction clearly, for example 11 + 12 = 23 or 23 − 11 = 12, so your method is visible.
- 💡Use the correct notation when writing nuclides, placing the mass number as a superscript before the symbol.
- 💡Write the symbol with the mass number clearly above and to the left, and the atomic number below and to the left.
- 💡Check that the numbers you use are consistent with the element's identity: the atomic number must match the number of protons.
- 💡When asked to identify an isotope, compare atomic numbers and mass numbers rather than just the element symbol.
- 💡When asked to define isotopes, include both the same element or same proton number and the different neutron number in one sentence.
- 💡For calculations, write the proton, neutron and electron counts separately before combining them, so an arithmetic slip is easy to spot.
- 💡If asked why isotopes react similarly, link the answer to the electron arrangement rather than to the nucleus.
- 💡Count protons and electrons separately before deciding the charge, then compare the two totals.
- 💡Use the group number of a main-group element to predict the charge of its ion, for example Group 1 gives 1⁺ and Group 2 gives 2⁺.
- 💡Check that the number of electrons lost matches the size of the positive charge, so the two parts of the answer agree.
- 💡When interpreting a representation such as ³⁷Cl⁻, state the number of protons, neutrons and electrons before explaining the charge and mass.
- 💡Use the periodic table to check the atomic number, then calculate neutrons as mass number minus atomic number.
- 💡For comparison questions, refer to both isotopes explicitly and link each difference to the representation, for example the mass number superscript or the charge superscript.
Common Mistakes
- Saying atoms contain equal numbers of protons and neutrons; correct this by stating the equal numbers are protons and electrons, while neutron number can differ.
- Treating a sodium ion as an atom; correct this by noting Na⁺ has lost one electron and so is charged, not neutral.
- Believing electrons contribute mass number; correct this by explaining mass number counts protons and neutrons only.
- Confusing atomic number with mass number; correct this by stating atomic number counts protons only, while mass number counts protons plus neutrons.
- Thinking isotopes of an element have different atomic numbers; correct this by noting isotopes differ in neutrons, not protons.
- Assuming atomic number equals neutron number; correct this by using an example such as carbon-14, which has six protons and eight neutrons.
- Including electrons in the mass number. Correction: only protons and neutrons are counted because electrons have negligible mass.
- Confusing mass number with atomic number. Correction: atomic number is the number of protons only, while mass number is protons plus neutrons.
- Writing the mass number as a subscript. Correction: mass number is written as a superscript before the symbol, for example ²³Na, while atomic number is written as a subscript.
- Swapping the positions of mass number and atomic number. Correction: mass number is the superscript and atomic number is the subscript.
- Thinking the atomic number equals the number of neutrons. Correction: atomic number equals the number of protons; neutrons are found by subtracting atomic number from mass number.
- Assuming all atoms of an element have the same mass number. Correction: isotopes have different mass numbers, so the mass number must be stated for a particular nuclide.
- Saying isotopes have different numbers of protons: correct this by stating that the proton number is fixed for an element and only the neutron number changes.
- Confusing mass number with relative atomic mass: correct this by describing mass number as a count of protons plus neutrons in one atom, while relative atomic mass is a weighted mean for the element.
- Writing an isotope as a charge, such as ³⁵Cl⁻: correct this by using mass number and atomic number notation, for example ³⁵₁₇Cl, and reserving charge symbols for ions.
- Saying that atoms become positive ions by gaining protons: correct this by stating that protons are fixed in the nucleus and it is the loss of electrons that creates the positive charge.
- Writing the charge after the symbol on the baseline, such as Na+: correct this by writing the charge as a superscript, Na⁺.
- Thinking that any electron can be lost equally easily: correct this by explaining that outer electrons are lost first because they are furthest from the nucleus and least strongly attracted.
- Confusing mass number with relative atomic mass. Correction: mass number is the total number of protons and neutrons in a specific isotope, while relative atomic mass is a weighted average for the element.
- Thinking that isotopes have different chemical identities because they have different masses. Correction: isotopes of the same element have the same atomic number and the same number of protons, so they are the same element.
- Writing the charge as a mass number or omitting the charge superscript for ions. Correction: use a separate superscript for charge, such as ³⁵Cl⁻, and a superscript for mass number, such as ³⁵Cl.