Group 7 — AQA GCSE Chemistry
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Group 7 explained
Group 7 covers 7 specification points, including The elements in Group 7 of the periodic table… and The halogens are non-metals and consist of….
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The elements in Group 7 of the periodic table are known as the halogens and have similar reactions because they all have seven electrons in their outer shell.
Group 7 contains fluorine, chlorine, bromine, iodine and astatine, collectively called the halogens. Each halogen atom has seven electrons in its outer shell, one short of a full shell. This shared outer-shell feature explains why halogens show similar reactions: they gain one electron to form a 1⁻ ion, or share one electron pair in a covalent bond, so completing the outer shell. Similarity of reactions does not mean identical reactivity. Down the group, atomic radius increases and shielding increases, so the nucleus attracts an incoming electron less strongly. Reactivity therefore decreases down Group 7: fluorine is the most reactive halogen and iodine is much less reactive. A displacement reaction, such as chlorine displacing bromine from potassium bromide solution, provides evidence for this trend.
The halogens are non-metals and consist of molecules made of pairs of atoms.
The halogens are non-metals, so they lack the typical properties of metals such as conduction of electricity and malleability. Their atoms have seven outer-shell electrons and need one more to complete the shell. Two halogen atoms can each contribute one electron to a shared pair, forming a single covalent bond. The result is a diatomic molecule, meaning a molecule made of two atoms, such as F₂, Cl₂, Br₂ and I₂. These small molecules are held together in the element by weak intermolecular forces, which explains why fluorine and chlorine are gases, bromine is a liquid and iodine is a solid at room temperature. The diatomic molecules are a key reason halogens are studied as pairs rather than single atoms.
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 to form a positive ion and the halogen forms a 1⁻ ion, so the compound is an ionic halide. For example, sodium reacts with chlorine to form sodium chloride, containing Na⁺ and Cl⁻ ions. When a halogen reacts with a non-metal such as hydrogen, electrons are shared rather than transferred, so the compound is covalent. Hydrogen and chlorine form hydrogen chloride, a covalent molecule containing a single covalent bond. The halogen always achieves a full outer shell, either by gaining an electron to form a 1⁻ ion or by sharing an electron pair. The nature of the compound therefore depends on whether the other element is a metal or a non-metal.
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. Their relative molecular mass increases as you descend from fluorine to iodine because each atom has more protons and neutrons. As the molecules get larger, the intermolecular forces between them become stronger. More energy is therefore needed to overcome these intermolecular forces, so melting point and boiling point rise steadily down the group. For example, fluorine and chlorine are gases at room temperature, bromine is a liquid, and iodine is a solid. The trend in melting and boiling points is explained entirely by the strength of the intermolecular forces, not by the covalent bonds within the molecules.
In Group 7, the reactivity of the elements decreases going down the group.
Halogens form 1⁻ ions in reactions with metals and in displacement reactions with metal halide solutions. Reactivity decreases down Group 7 from fluorine to iodine because the atoms become larger. As atomic radius increases, the outer shell is further from the nucleus and experiences greater shielding from inner electron shells. Consequently, the electrostatic attraction between the nucleus and an incoming electron becomes weaker, making it harder to gain an electron. For example, chlorine displaces bromine from aqueous potassium bromide (Cl₂ + 2KBr → 2KCl + Br₂) because chlorine is more reactive and gains an electron more easily than bromine.
A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt.
A displacement reaction occurs when a more reactive halogen is added to an aqueous solution containing a less reactive halide salt. The more reactive halogen gains an electron and forms a 1⁻ ion, while the less reactive halide ion loses an electron and becomes a halogen molecule. For example, chlorine displaces bromine from potassium bromide solution because chlorine is above bromine in Group 7. The colour change observed depends on the halogen formed: bromine gives an orange solution and iodine gives a brown solution. This reaction demonstrates the reactivity trend and can be used to place halogens in order of reactivity.
Students should be able to: explain how properties of the elements in Group 7 depend on the outer shell of electrons of the atoms predict properties from given trends down the group.
The properties of Group 7 elements are determined by their outer shell electron configuration. Each halogen atom has seven electrons in its outer shell, so it needs to gain one electron to achieve a stable noble gas configuration. This explains why halogens form 1⁻ ions and why reactivity depends on how strongly the nucleus attracts an incoming electron. Down the group, atomic radius increases and shielding increases, so the attraction weakens and reactivity decreases. Students should use these ideas to predict properties such as physical state, relative melting point and displacement behaviour from trends, rather than memorising isolated facts.
Your focus
- State the outer-shell electron configuration shared by all halogen atoms.
- Explain how gaining one electron or sharing an electron pair leads to similar halogen reactions.
- Describe how reactivity changes down Group 7 and relate it to atomic size and shielding.
Show all 21 objectives
- Identify the halogens as non-metals that exist as diatomic molecules.
- Describe the single covalent bond in a halogen molecule as a shared pair of electrons.
- Relate weak intermolecular forces to the physical state of a halogen at room temperature.
- 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 the halogen completes its outer shell by gaining an electron or sharing an electron pair.
- Describe the trend in relative molecular mass, melting point and boiling point down Group 7.
- Explain the trend in terms of intermolecular forces between diatomic molecules.
- Use the trend to predict the physical state or relative melting point of a halogen from its position in the group.
- State that reactivity decreases down Group 7.
- Explain the decrease in reactivity in terms of atomic radius, shielding and attraction for an incoming electron.
- Relate the reactivity trend to displacement reactions between halogens and halide salts.
- Describe displacement reactions between halogens and aqueous halide salts.
- Predict whether a displacement reaction will occur using the Group 7 reactivity trend.
- Interpret colour changes to identify the halogen produced in a displacement reaction.
- Explain how the outer shell electron configuration of Group 7 atoms determines their properties.
- Use trends down Group 7 to predict physical and chemical properties of halogens.
- Link reactivity, ion formation and displacement behaviour to the outer shell electrons of halogen atoms.
Group 7 exam tips
Marking Points
- Identify the halogens as the elements in Group 7 of the periodic table.
- State that each halogen atom has seven electrons in its outer shell.
- Explain that halogens gain one electron to form a 1⁻ ion or share one electron pair in a covalent bond.
- Link the common outer-shell configuration to the similar reactions of the halogens.
- Explain that reactivity decreases down Group 7 because the outer shell is further from the nucleus and more shielded.
- Use a displacement reaction as evidence that a more reactive halogen can displace a less reactive one from its compound.
- Classify the halogens as non-metals.
- State that halogen molecules are diatomic, containing two atoms joined by a single covalent bond.
- Write the molecular formulae of common halogens as F₂, Cl₂, Br₂ and I₂.
- Explain that each atom contributes one electron to a shared pair, completing the outer shell of both atoms.
- Relate weak intermolecular forces between diatomic molecules to the physical states of the halogens at room temperature.
- State that halogens react with metals to form ionic compounds called halides.
- Describe the metal atom losing electrons and the halogen atom gaining one electron to form a 1⁻ ion.
- Give an example such as sodium chloride, formed from Na⁺ and Cl⁻ ions.
- State that halogens react with non-metals such as hydrogen to form covalent compounds.
- Describe the covalent bond in hydrogen chloride as a shared pair of electrons.
- Explain that the halogen completes its outer shell either by gaining one electron or by sharing an electron pair.
- Group 7 elements are diatomic molecules, so relative molecular mass increases down the group as atomic number increases.
- Larger molecules have stronger intermolecular forces between them.
- Stronger intermolecular forces require more energy to overcome, so melting point and boiling point increase down the group.
- The trend is a physical property trend explained by intermolecular forces, not by the strength of the covalent bonds within each molecule.
- State symbols at room temperature change from gas to liquid to solid: fluorine and chlorine are gases, bromine is a liquid, iodine is a solid.
- Halogens react by gaining one electron to form a 1⁻ ion in reactions with metals and in displacement reactions.
- Down the group, atomic radius increases and there is more shielding.
- The increased distance and shielding result in a weaker electrostatic attraction between the nucleus and an incoming electron.
- Fluorine is the most reactive halogen and iodine is the least reactive of the commonly studied halogens.
- The reactivity trend is linked to the outer shell electron configuration and the ease of gaining an electron.
- A more reactive halogen displaces a less reactive halogen from an aqueous solution of its salt.
- The more reactive halogen gains an electron and forms a 1⁻ ion, while the less reactive halide ion is oxidised to form a halogen molecule.
- Chlorine displaces bromine from potassium bromide solution and iodine from potassium iodide solution.
- Bromine displaces iodine from potassium iodide solution but cannot displace chlorine from potassium chloride solution.
- The colour of the solution changes as the halogen formed appears, for example orange for bromine and brown for iodine.
- Group 7 atoms have seven electrons in their outer shell and gain one electron to form a 1⁻ ion.
- The ease of gaining an electron depends on the attraction between the nucleus and the incoming electron, which is affected by atomic radius and shielding.
- Down the group, atomic radius and shielding increase, so the attraction for an incoming electron decreases and reactivity decreases.
- Physical properties such as melting point and boiling point increase down the group because intermolecular forces become stronger.
- Predictions should be based on trends in the group, such as a halogen lower down being less reactive and having a higher melting point.
Examiner Tips
- 💡Begin similarity answers with the outer-shell electron count, because that is the cause of the shared chemistry.
- 💡When explaining reactivity trends, compare distance from the nucleus and shielding between the two named halogens.
- 💡For displacement questions, name the more reactive halogen and state which ion it replaces.
- 💡Use superscript charges for all ions, for example F⁻, Cl⁻, Br⁻ and I⁻.
- 💡When naming a halogen molecule, always include the subscript two, for example Br₂.
- 💡If asked why halogens form diatomic molecules, refer to the need to complete the outer shell by sharing one electron pair.
- 💡Link physical state to intermolecular forces rather than to the strength of the covalent bond.
- 💡Check that formulae use subscript numerals, not baseline numerals.
- 💡Decide first whether the other element is a metal or a non-metal, because that determines ionic or covalent bonding.
- 💡Name the compound carefully, using the halide ending for ionic compounds, for example sodium bromide.
- 💡Show electron transfer for metals and electron sharing for non-metals in your explanation.
- 💡Use superscript charges for ions and subscript numerals for formulae, for example MgCl₂.
- 💡Link each property to the structure and bonding explanation: larger molecules, stronger intermolecular forces, more energy needed to melt or boil.
- 💡Use the correct comparison language, such as higher, lower, stronger and weaker, rather than vague words like bigger or more.
- 💡When asked to predict a trend, refer to the position in Group 7 and the change in intermolecular forces rather than memorising colours alone.
- 💡Explain reactivity using the outer shell electron and the attraction from the nucleus, not just by stating the trend.
- 💡Use comparative terms such as more reactive and less reactive, and name the halogens correctly.
- 💡Identify which halogen is more reactive by its position in Group 7 before predicting whether displacement occurs.
- 💡Describe the colour change and name the halogen responsible for it.
- 💡Write a word equation or balanced symbol equation to show the displacement clearly.
- 💡Start explanations with the outer shell electron configuration and then link to the property being explained.
- 💡When predicting a property, state the trend and use it to justify the prediction.
- 💡Use correct chemical terminology, including 1⁻ ions, atomic radius, shielding and intermolecular forces.
Common Mistakes
- Saying halogens have seven electrons in total rather than seven in the outer shell; correct this by specifying the outer shell.
- Writing halide ions with a baseline numeral and postfix sign; correct this by using superscript forms such as Cl⁻ and Br⁻.
- Claiming all halogens react identically; correct this by saying reactions are similar in type but reactivity decreases down the group.
- Confusing halogens with halide ions; correct this by noting that halogens are the elements, while halides are the 1⁻ ions formed from them.
- Writing halogen molecules as single atoms, for example Cl; correct this by writing Cl₂ to show the diatomic molecule.
- Describing the bond within a halogen molecule as ionic; correct this by describing a single covalent bond formed from a shared pair of electrons.
- Confusing intermolecular forces with covalent bonds; correct this by stating that weak forces act between molecules while strong covalent bonds act within each molecule.
- Using baseline numerals for formulae; correct this by using subscript numerals such as F₂ and I₂.
- Describing hydrogen chloride as ionic; correct this by describing it as a covalent molecule with a shared pair of electrons.
- Writing halide ions with baseline numerals and postfix signs; correct this by using superscript forms such as Cl⁻ and Br⁻.
- Saying the halogen loses an electron when reacting with a metal; correct this by stating that the halogen gains one electron.
- Confusing the element name chlorine with the ion name chloride; correct this by using chloride for the 1⁻ ion in compounds.
- Saying that melting point increases because the covalent bonds get stronger; the correction is that intermolecular forces between molecules get stronger, while covalent bonds within molecules do not break during melting.
- Confusing relative molecular mass with relative atomic mass; the correction is that halogens are diatomic, so relative molecular mass is twice the relative atomic mass for each element.
- Claiming that all halogens are gases at room temperature; the correction is that bromine is a liquid and iodine is a solid under standard conditions.
- Saying reactivity decreases because the atoms get heavier; the correction is that reactivity depends on the attraction for an incoming electron, not on mass.
- Claiming that increased nuclear charge always increases reactivity; the correction is that increased distance and shielding outweigh the extra nuclear charge down the group.
- Stating that halogens only form 1⁻ ions when reacting with metals; the correction is that they also form 1⁻ ions in displacement reactions with metal halide solutions.
- Writing that a less reactive halogen displaces a more reactive one; the correction is that the more reactive halogen displaces the less reactive halogen.
- Forgetting that the halide must be in aqueous solution; the correction is that displacement is described for an aqueous solution of the salt.
- Confusing the colour of the halogen with the colour of the halide ion; the correction is that the colour change is due to the halogen molecule formed, not the halide ion.
- Saying halogens lose seven electrons to form ions; the correction is that they gain one electron to form a 1⁻ ion.
- Treating each property as an isolated fact rather than linking it to the outer shell electron configuration and atomic size.
- Predicting that reactivity increases down the group because atomic number increases; the correction is that reactivity decreases because the outer shell electron is further from the nucleus and more shielded.