Group 7 — AQA GCSE Combined Science
Test yourself on Group 7 with AQA GCSE practice questions.
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Group 7 explained
Group 7 elements, the halogens, sit in the second-from-right column of the periodic table.
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Each halogen atom has seven electrons in its outer shell, so all halogens react in similar ways: they gain one electron to complete a stable outer shell of eight. This shared outer-shell pattern explains why fluorine, chlorine, bromine, iodine and astatine behave as a family. Halogens are non-metals, so they are poor conductors and have low melting and boiling points compared with metals. Their atoms pair up covalently, forming diatomic molecules such as F₂, Cl₂, Br₂ and I₂. For example, chlorine exists as Cl₂, not as single Cl atoms, because the two atoms share a pair of electrons. The seven outer electrons and diatomic molecules together explain the similar reactions and the physical properties of the group.
Students should be able to describe the nature of the compounds formed when chlorine, bromine and iodine react with metals and non-metals.
When chlorine, bromine or iodine reacts with a metal, the halogen gains one electron from the metal atom. The metal loses electrons and becomes a positive ion, while the halogen becomes a negative ion. The oppositely charged ions attract strongly, forming an ionic compound. For example, sodium reacts with chlorine to form sodium chloride, NaCl, containing Na⁺ and Cl⁻ ions. When a halogen reacts with a non-metal such as hydrogen, the atoms share electrons instead. Hydrogen and chlorine share a pair of electrons to form hydrogen chloride, HCl, a simple covalent molecule. So the nature of the compound depends on the other element: ionic with metals, covalent with non-metals. In both cases the halogen achieves a full outer shell of eight electrons, which is why it reacts.
In Group 7, the further down the group an element is the higher its relative molecular mass, melting point and boiling point.
Group 7 halogens exist as diatomic molecules, so their relative molecular mass increases as you go down the group because each atom has more protons and neutrons. For example, fluorine is F₂ (Mr = 38), chlorine is Cl₂ (Mr = 71) and bromine is Br₂ (Mr = 160). Larger molecules have more electrons and stronger London dispersion forces between molecules, so more energy is needed to overcome these intermolecular forces. This explains why melting point and boiling point increase down the group: fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid. The trend is assessed by describing the state and trend down the group and linking it to intermolecular forces, not by memorising exact temperatures.
In Group 7, the reactivity of the elements decreases going down the group.
Group 7 halogens react by gaining one electron to form a 1⁻ ion. Going down the group, the outer shell becomes further from the nucleus and there are more inner shells shielding the outer electrons. Although nuclear charge increases, the increased distance and shielding mean the attraction for an incoming electron is weaker, so reactivity decreases. This is shown by displacement reactions: a more reactive halogen higher in the group displaces a less reactive halide lower down. For example, chlorine displaces bromine from potassium bromide solution, but bromine cannot displace chlorine from potassium chloride solution. The trend is assessed by explaining electron gain and by predicting displacement reactions.
A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt.
Group 7 halogens gain one electron to form 1⁻ ions, and reactivity decreases down the group as atomic radius increases and shielding rises, so the outer shell is gained less easily. A more reactive halogen therefore oxidises a less reactive halide ion in aqueous solution, taking its place in the salt. For example, chlorine added to potassium bromide solution displaces bromine: Cl₂ + 2KBr → 2KCl + Br₂. The solution turns orange-brown as bromine forms. Iodine is displaced by chlorine or bromine, giving a brown solution. Fluorine is too hazardous for school tests. Displacement is evidence for the reactivity trend, and the colour change identifies which halogen has formed.
explain how properties of the elements in Group 7 depend on the outer shell of electrons of the atoms
Group 7 elements are the halogens: fluorine, chlorine, bromine, iodine and astatine. Each atom has seven electrons in its outer shell, giving it a configuration that is one electron short of a stable noble-gas arrangement. This single feature explains most of their chemistry. When a halogen atom gains one electron it forms a 1⁻ ion with a full outer shell, so halogens are reactive non-metals and strong oxidising agents. Two halogen atoms can also share a pair of electrons in a single covalent bond, forming diatomic molecules such as Cl₂. Because the outer shell is one electron short, the halogens react vigorously with metals such as sodium, forming salts like sodium chloride. Reactivity decreases down the group because the outer shell is further from the nucleus and shielded by more inner shells, so the attraction for an incoming electron is weaker.
predict properties from given trends down the group.
Down Group 7, the halogens show clear trends in physical and chemical properties. Fluorine and chlorine are gases, bromine is a liquid and iodine is a solid, so melting and boiling points increase down the group. The colour deepens from pale yellow through green, red-brown and grey-black. Reactivity decreases down the group because the outer shell is further from the nucleus and shielded by more inner shells, so the attraction for an incoming electron is weaker. These trends allow predictions about less familiar halogens such as astatine. Astatine should be a dark solid with a high melting point and low reactivity, and it should form a 1⁻ ion and diatomic molecules. Predictions are made by identifying the trend, checking it across at least two known elements and extending it to the unknown element.
Your focus
- State that Group 7 elements are the halogens and identify their position in the periodic table.
- Explain that halogens react similarly because each atom has seven outer-shell electrons.
- Describe halogens as non-metals that form diatomic molecules such as Cl₂ and Br₂.
Show all 21 objectives
- Describe the ionic compounds formed when chlorine, bromine or iodine reacts with a metal.
- Describe the covalent compounds formed when a halogen reacts with a non-metal such as hydrogen.
- Explain how electron transfer or sharing allows the halogen atom to complete its outer shell.
- State that Group 7 elements exist as diatomic molecules and that relative molecular mass increases down the group.
- Explain that increased molecular size leads to stronger London dispersion forces between molecules.
- Relate stronger intermolecular forces to higher melting points and boiling points and to the observed states at room temperature.
- Describe that Group 7 reactivity decreases down the group.
- Explain the trend in terms of electron gain, atomic radius and shielding.
- Use the trend to predict and interpret displacement reactions between halogens and halide ions.
- Describe displacement reactions between halogens and aqueous solutions of halide salts.
- Explain displacement in terms of the reactivity trend down Group 7.
- Write balanced equations for halogen displacement reactions and identify the halogen formed from colour changes.
- Describe the outer-shell electron configuration of Group 7 atoms and the 1⁻ ion they form.
- Explain how the outer shell accounts for the reactivity and bonding of halogens.
- Relate the decrease in reactivity down Group 7 to increasing distance and shielding of the outer shell.
- Describe the trends in melting point, colour and reactivity down Group 7.
- Use trends from known halogens to predict properties of less familiar halogens such as astatine.
- Justify predictions by referring to outer-shell electrons and distance from the nucleus.
Group 7 exam tips
Marking Points
- Group 7 elements are called halogens and include fluorine, chlorine, bromine, iodine and astatine.
- Each halogen atom has seven electrons in its outer shell, giving the group its similar chemical reactions.
- Halogens are non-metals, so they are poor conductors of electricity and have relatively low melting and boiling points.
- Halogen atoms form diatomic molecules by sharing a pair of electrons, for example Cl₂, Br₂ and I₂.
- The tendency to gain one electron to complete an outer shell of eight explains the similar reactivity of the halogens.
- Halogens react with metals to form ionic compounds in which the halogen becomes a negative ion, for example Cl⁻.
- The metal atom loses electrons to become a positive ion, for example Na⁺, and the ions attract to form the compound.
- Halogens react with non-metals such as hydrogen to form covalent compounds in which electrons are shared.
- Hydrogen chloride, HCl, is a covalent molecule formed when hydrogen and chlorine share a pair of electrons.
- In both ionic and covalent compounds, the halogen atom gains or shares electrons to complete an outer shell of eight.
- Halogens are diatomic molecules, so relative molecular mass increases down the group as atoms gain more protons and neutrons.
- Larger molecules have more electrons, leading to stronger London dispersion forces between molecules.
- Stronger intermolecular forces require more energy to overcome, so melting point and boiling point increase down the group.
- The physical states at room temperature change from gas (F₂, Cl₂) to liquid (Br₂) to solid (I₂), consistent with increasing melting and boiling points.
- The trend applies to the elements in Group 7 and is explained by intermolecular forces, not by covalent bond strength within each molecule.
- Halogens react by gaining one electron to form a 1⁻ ion, so reactivity depends on how easily the outer shell attracts an electron.
- Down the group, the outer shell is further from the nucleus and there are more inner shells, increasing shielding.
- The increased distance and shielding outweigh the increased nuclear charge, so the attraction for an incoming electron is weaker.
- A more reactive halogen higher in the group displaces a less reactive halide lower in the group from its compound.
- Displacement observations, such as a colour change in solution, provide evidence for the reactivity trend.
- State that halogens react by gaining one electron to form 1⁻ ions, so a more reactive halogen is a stronger oxidising agent.
- Explain that reactivity decreases down Group 7 because atomic radius and shielding increase, so the outer shell attracts an electron less strongly.
- Write a balanced displacement equation, such as Cl₂ + 2KBr → 2KCl + Br₂, keeping formulae and state symbols correct.
- Describe the observation: a colour change in the aqueous solution as the displaced halogen forms, for example colourless to orange-brown with bromine.
- Identify a displacement reaction from given results by comparing the halogen added with the halogen in the salt.
- Explain that no reaction occurs when the added halogen is less reactive than the halogen in the salt, for example iodine added to potassium chloride solution.
- State that a Group 7 atom has seven electrons in its outer shell.
- Explain that gaining one electron gives a full outer shell and forms a 1⁻ ion, for example Cl + e⁻ → Cl⁻.
- Link the tendency to gain an electron to reactivity as non-metals and oxidising agents.
- Explain that two atoms can share electrons to form diatomic molecules such as Cl₂.
- Describe the trend in reactivity down the group in terms of increased distance and shielding of the outer shell from the nucleus.
- Use a named example, such as chlorine reacting with sodium to form sodium chloride, to support the explanation.
- Identify the trend being used, such as increasing melting point or decreasing reactivity down Group 7.
- Check the trend against at least two known halogens before extending it to the unknown element.
- Predict a physical property, for example that astatine is a solid with a higher melting point than iodine.
- Predict a chemical property, for example that astatine is less reactive than iodine and forms a 1⁻ ion.
- Justify the prediction using the outer-shell electron configuration and increasing distance and shielding from the nucleus.
- State that predictions from trends are tentative and may need checking against data.
Examiner Tips
- 💡Link the seven outer electrons directly to the similar reactions, rather than listing the two facts separately.
- 💡Use the correct diatomic formula with a subscript, for example Cl₂, when describing a halogen molecule.
- 💡Contrast halogen properties with metal properties, such as poor conduction and low melting points, to show understanding of non-metals.
- 💡Name the type of bonding first, then explain the electron transfer or sharing that produces it.
- 💡Use correct ion symbols with superscript charges, such as Na⁺ and Cl⁻, when describing ionic compounds.
- 💡Give one metal example and one non-metal example, such as sodium chloride and hydrogen chloride, to cover both parts of the statement.
- 💡When describing the trend, name the halogens in order and state the change in state from gas to liquid to solid.
- 💡Use the phrase London dispersion forces or intermolecular forces, and link stronger forces to more energy needed for melting or boiling.
- 💡If asked to explain, compare two named halogens, such as chlorine and bromine, to show the link between Mr, forces and boiling point.
- 💡Link reactivity to electron gain and the formation of a 1⁻ ion, then explain the effect of distance and shielding.
- 💡For displacement questions, name the halogen and the halide, state which is more reactive, and describe the expected observation.
- 💡Use comparative language such as more reactive than and less reactive than to make the trend clear.
- 💡Link every observation to the identity of the halogen formed, for example an orange-brown colour indicates bromine.
- 💡Balance displacement equations by counting halogen atoms on both sides before checking potassium and other spectator ions.
- 💡When comparing two halogens, name the more reactive one first and state clearly which halide it displaces.
- 💡Always link the outer-shell electron count to a specific consequence, such as ion formation, covalent bonding or reactivity, rather than just stating the number seven.
- 💡When explaining the reactivity trend, refer explicitly to distance from the nucleus and shielding by inner electron shells.
- 💡Use correct notation for ions and molecules, for example Cl⁻ and Cl₂, and balance any equation you write.
- 💡Name the trend you are using and quote data for at least two known halogens to support your prediction.
- 💡Separate physical predictions, such as state and colour, from chemical predictions, such as reactivity and ion formed.
- 💡Use correct notation for ions and molecules, for example I⁻ and I₂, and avoid vague words such as better or worse without explanation.
Common Mistakes
- Writing halogen molecules as single atoms such as Cl or Br; correct this by writing the diatomic forms Cl₂, Br₂ and I₂.
- Saying halogens have eight outer electrons; correct this by stating they have seven outer electrons and gain one electron to reach eight.
- Describing halogens as metals because they react with metals; correct this by stating halogens are non-metals that form compounds with metals.
- Saying halogens form covalent compounds with all elements; correct this by stating they form ionic compounds with metals and covalent compounds with non-metals.
- Writing sodium chloride as molecules of NaCl rather than as a lattice of Na⁺ and Cl⁻ ions; correct this by describing the ionic lattice.
- Forgetting that the halogen becomes a negative ion in an ionic compound; correct this by showing Cl gaining one electron to form Cl⁻.
- Thinking that relative molecular mass increases because atoms get bigger without linking it to more protons and neutrons; correct by stating that each successive halogen atom has more nucleons, so the diatomic molecule has greater Mr.
- Confusing intermolecular forces with covalent bonds; correct by explaining that London dispersion forces between molecules are overcome during melting and boiling, not the covalent bonds inside each molecule.
- Assuming all halogens are gases at room temperature; correct by recalling that bromine is a liquid and iodine is a solid, which supports the increasing melting and boiling point trend.
- Saying reactivity increases down the group because atoms get larger; correct by explaining that larger atoms have weaker attraction for an incoming electron, so reactivity decreases.
- Ignoring shielding and only referring to distance; correct by stating that more inner shells shield the outer shell and reduce the effective nuclear attraction.
- Predicting that a lower halogen can displace a higher halide; correct by applying the rule that the more reactive halogen higher in the group does the displacing.
- Writing that the more reactive halogen is displaced, rather than the less reactive halogen: correct this by saying the more reactive halogen displaces the less reactive halogen from its salt.
- Using the formula of the halogen salt incorrectly, such as KBr₂ instead of KBr: correct this by checking that potassium forms 1⁺ ions and bromide forms 1⁻ ions, giving KBr.
- Claiming that reactivity increases down Group 7: correct this by linking decreasing reactivity to increasing atomic radius and shielding, so the outer shell gains an electron less easily.
- Saying halogens lose seven electrons to gain a full outer shell; correct this by stating that gaining one electron is far more favourable and forms a 1⁻ ion.
- Writing that reactivity increases down Group 7; correct this by explaining that the outer shell is further from the nucleus and more shielded, so the attraction for an incoming electron is weaker.
- Confusing the outer-shell electron count with the group number for other groups; correct this by stating that for Group 7 the group number equals the number of outer-shell electrons, which is seven.
- Assuming all halogens are gases; correct this by recalling that bromine is a liquid and iodine is a solid at room temperature.
- Predicting that reactivity increases down the group; correct this by linking decreasing reactivity to weaker attraction for an incoming electron as the outer shell becomes further from the nucleus.
- Extending a trend from only one known element; correct this by checking the trend across at least two known halogens before predicting.