The periodic table — AQA GCSE Combined Science
Test yourself on The periodic table with AQA GCSE practice questions.
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The periodic table explained
The modern periodic table orders elements by increasing atomic (proton) number, the number of protons in the nucleus.
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This ordering places elements with similar chemical properties into vertical columns called groups. For example, Group 1 contains lithium, sodium and potassium, which all react vigorously with water and form 1⁺ ions. The arrangement is called periodic because similar properties recur at regular intervals as proton number increases. Reading across a period, properties change gradually, but moving down a group, behaviour repeats in a pattern. This structure allows chemists to predict the properties and reactions of an element from its position, rather than learning each element separately.
Elements in the same group in the periodic table have the same number of electrons in their outer shell (outer electrons) and this gives them similar chemical properties.
Elements in the same group have identical numbers of outer-shell electrons. For example, Group 1 elements each have one outer electron, and Group 7 elements each have seven. Chemical properties depend mainly on outer electrons because reactions involve losing, gaining or sharing these electrons to achieve a stable electron arrangement. Sodium (2,8,1) and potassium (2,8,8,1) both have one outer electron, so both form 1⁺ ions and react similarly with water, even though potassium reacts more vigorously. The shared outer-electron count explains why group members behave alike, while the increasing number of inner shells down a group explains trends in reactivity.
explain how the position of an element in the periodic table is related to the arrangement of electrons in its atoms and hence to its atomic number
The periodic table is arranged in order of increasing atomic number, and this order reflects the number of protons and electrons in each atom. Elements in the same group have the same number of outer-shell electrons, which is why they show similar chemical properties. Elements in the same period have the same number of occupied electron shells. For example, chlorine has atomic number 17, so it has 17 protons and 17 electrons, with arrangement 2,8,7; it is in Group 7 and Period 3, so it has seven outer electrons and three occupied shells. You should be able to move between atomic number, electron arrangement and periodic table position in either direction, and use this to explain similarities within groups and trends across periods.
predict possible reactions and probable reactivity of elements from their positions in the periodic table.
The periodic table arranges elements by increasing atomic number, so elements in the same group have the same number of outer-shell electrons and therefore similar chemical properties. To predict reactivity, first locate the element's group and period. Group 1 metals become more reactive down the group because the outer electron is lost more easily as atomic radius increases and shielding increases. Group 7 non-metals become less reactive down the group because it becomes harder to gain an electron. Group 0 noble gases are unreactive due to full outer shells. For example, potassium is below sodium in Group 1, so it reacts more vigorously with water, producing hydrogen and potassium hydroxide. You can also predict likely reaction partners: Group 1 metals react with Group 7 non-metals to form ionic salts such as sodium chloride.
Your focus
- State that the modern periodic table is ordered by increasing atomic (proton) number.
- Identify groups as vertical columns containing elements with similar properties.
- Explain the meaning of periodic in terms of regularly repeating properties.
Show all 12 objectives
- Identify the number of outer-shell electrons for elements in a given main group.
- Explain how a shared outer-electron count leads to similar chemical properties.
- Use electron configurations to compare elements in the same group.
- Explain the link between atomic number and the number of electrons in a neutral atom.
- Use electron arrangement to justify the group and period of an element.
- Explain why elements in the same group have similar chemical properties.
- Locate an element's group and period and use these to state its outer-shell electron arrangement.
- Describe and apply the reactivity trends down Group 1 and Group 7 and explain the lack of reactivity of Group 0.
- Predict possible reactions and probable reactivity of named elements from their positions in the periodic table, giving a reason based on electron configuration.
The periodic table exam tips
Marking Points
- States that elements are arranged in order of increasing atomic (proton) number.
- Identifies vertical columns as groups and links them to similar chemical properties.
- Explains that the table is periodic because similar properties recur at regular intervals.
- Uses a named group, such as Group 1 or Group 7, to illustrate similar properties.
- Distinguishes atomic (proton) number from mass number when describing the ordering.
- States that elements in the same group have the same number of outer-shell electrons.
- Links outer-shell electron number to similar chemical properties.
- Gives a correct example, such as Group 1 elements having one outer electron or Group 7 elements having seven.
- Explains that reactions involve outer electrons being lost, gained or shared.
- Uses electron configurations to show the shared outer-electron count, for example sodium 2,8,1 and potassium 2,8,8,1.
- State that elements are arranged in order of increasing atomic number.
- Link atomic number to the number of protons and electrons in a neutral atom.
- Explain that elements in the same group have the same number of outer-shell electrons.
- Explain that elements in the same period have the same number of occupied electron shells.
- Use electron arrangement to explain why elements in a group have similar chemical properties.
- Deduce the position of an element from a given electron arrangement or atomic number.
- Identify the element's group and period from its position in the periodic table.
- Relate group number to the number of outer-shell electrons and hence to similarity in chemical properties.
- State and apply the trend in reactivity down Group 1: reactivity increases as the outer electron is lost more easily.
- State and apply the trend in reactivity down Group 7: reactivity decreases as it becomes harder to gain an electron.
- Recognise that Group 0 elements are unreactive because they have full outer electron shells.
- Predict a plausible reaction by combining a reactive metal from Group 1 or 2 with a non-metal from Group 6 or 7, for example sodium with chlorine forming sodium chloride.
- Use a named example, such as potassium reacting more vigorously than sodium with water, to support a prediction.
Examiner Tips
- 💡Quote atomic (proton) number explicitly rather than saying atomic number alone if the question asks for the modern ordering principle.
- 💡When asked why the table is periodic, link the word to a repeating pattern of properties down a group.
- 💡Use a specific group example to support your explanation, such as the reactivity trend in Group 1.
- 💡Write electron configurations such as 2,8,1 to make the outer-electron count visible in your answer.
- 💡When explaining similar properties, always connect outer electrons to how elements lose, gain or share electrons in reactions.
- 💡If asked about a trend down a group, separate the similar outer-electron count from the changing reactivity.
- 💡When explaining a similarity, name the group and state the number of outer electrons before linking to reactivity.
- 💡If asked to explain the position of an element, work from atomic number to electron arrangement to group and period in a clear sequence.
- 💡Check that your electron arrangement uses commas between shell numbers and does not exceed 2, 8, 8 for the first three shells.
- 💡Always name the group and period of the element before making a prediction, so your reasoning is clear.
- 💡Link reactivity to electron loss or gain and mention the outer-shell electron arrangement.
- 💡Use a specific example, such as caesium reacting more vigorously than lithium with water, to show understanding of the trend.
- 💡If asked to predict a reaction, name the likely products and the type of reaction, for example a metal plus a non-metal forming an ionic compound.
Common Mistakes
- Saying elements are arranged in order of relative atomic mass: correct this by stating that the modern table uses atomic (proton) number, which resolves earlier inconsistencies.
- Confusing groups with periods: correct this by noting that groups are vertical columns and periods are horizontal rows.
- Describing the table as periodic because it repeats every eight elements only: correct this by explaining that similar properties recur at regular intervals, not necessarily every eight.
- Saying elements in the same group have the same total number of electrons: correct this by specifying outer-shell electrons only.
- Claiming all group members react with identical vigour: correct this by explaining that reactivity changes down a group even though the type of reaction is similar.
- Confusing group number with outer-electron number for every group: correct this by noting the link works for main-group elements, such as Group 1 and Group 7, but transition metals do not follow it simply.
- Saying elements are arranged by relative atomic mass: the error is ignoring the modern basis of the table; the correction is that the modern periodic table is ordered by increasing atomic number.
- Assuming all atoms of an element have the same number of neutrons: the error is confusing atomic number with mass number; the correction is that atomic number gives protons, while isotopes differ in neutrons.
- Thinking group number always equals outer electrons for every element: the error is applying the rule to transition metals; the correction is that the simple group-number rule applies to main-group elements.
- Assuming all Group 1 metals react identically: correct this by stating that reactivity increases down the group, so potassium reacts more vigorously than lithium.
- Confusing the trends in Group 1 and Group 7: correct this by remembering that Group 1 reactivity increases down the group, while Group 7 reactivity decreases down the group.
- Thinking noble gases are reactive because they are in the periodic table: correct this by explaining that Group 0 elements have full outer shells and are chemically unreactive.
- Predicting reactivity from atomic mass alone: correct this by using group position and outer-shell electron configuration as the main basis for prediction.