Ch 2 · The chemical properties of the halogen elements and the hydrogen halides · AS & A Level Chemistry 9701
Topic 2 of 4 in Group 17 — AS & A Level Chemistry 9701 — 48 lessons.
Introduction
The halogens are oxidising agents, and their oxidising power falls steadily down the group. This single trend explains why chlorine reacts vigorously with hydrogen while iodine barely does, why the hydrogen halides differ so much in thermal stability, and why one halogen displaces another from solution — all questions Cambridge asks in this subtopic.
Key ideas — Halogens as oxidising agents
Each halogen atom has seven outer-shell electrons and reacts by gaining one electron to form a stable, singly-charged halide ion with a noble-gas electron configuration: X_2 + 2e^- arrow 2X^-. Oxidising strength (the tendency to gain electrons) decreases down the group from Cl_2 to I_2, because atomic radius increases down the group, so an incoming electron is added further from the nucleus and experiences greater shielding from inner shells; the attraction between the nucleus and the incoming electron therefore weakens down the group.
🧑🏫 Teacher's explanation
Think of halogens as 'electron thieves'. They are very hungry for one more electron to complete their outer shell and reach a stable, noble-gas configuration. Because they are so hungry, they act as oxidising agents—they take electrons from other substances, causing those substances to be oxidised.
📖 Definitions
🔍 Worked example — Easy
Predict the products of the reaction between chlorine gas and potassium bromide solution.
🔍 Worked example — Standard
Explain why Cl_2 is a stronger oxidising agent than I_2.
🔍 Worked example — Exam-style
A student reacts Br_2(aq) with NaI(aq). State whether a reaction occurs and provide the balanced equation if it does.
⚠️ Common errors
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🎯 Exam question
Key ideas — Reactions with hydrogen
All halogens react with hydrogen to form a hydrogen halide, H_2(g) + X_2 arrow 2HX(g), but reactivity decreases sharply down the group. F_2 and Cl_2 react explosively or rapidly even in the dark or at room temperature (chlorine reacts explosively with hydrogen in direct sunlight); Br_2 reacts more slowly and needs a heated catalyst; I_2 reacts only slowly and reversibly even when heated with a catalyst, giving an equilibrium mixture: H_2(g) + I_2(g) leftharpoons 2HI(g). This decreasing reactivity mirrors the decreasing oxidising strength of the halogens down the group.
🧑🏫 Teacher's explanation
Think of the halogens as 'greedy' atoms looking to steal electrons to complete their outer shell. Because they are so hungry for electrons, they react with hydrogen to form stable compounds. However, as you go down Group 17, the atoms get much larger. This means the nucleus is further away from the incoming electrons, making the 'pull' or the 'greediness' (oxidising power) much weaker.
📖 Definitions
🔍 Worked example — Easy
Predict the state of the reaction between H_2(g) and Cl_2(g) at room temperature.
🔍 Worked example — Standard
Write the balanced equation for the reaction between H_2(g) and I_2(g) and indicate if it is reversible.
🔍 Worked example — Exam-style
Calculate the enthalpy change Δ H for the reaction H_2(g) + Cl_2(g) arrow 2HCl(g) if the bond enthalpy of H-Cl is 431 kJ/mol and H-H and Cl-Cl are 436 kJ/mol and 243 kJ/mol respectively.
⚠️ Common errors
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🎯 Exam question
Key ideas — Thermal stability of the hydrogen halides
Thermal stability of the hydrogen halides decreases down the group: HF and HCl are stable to heat, HBr decomposes only slightly on strong heating, and HI decomposes noticeably even at moderate temperatures, 2HI(g) leftharpoons H_2(g) + I_2(g). This trend is explained directly by H–X bond enthalpy, which falls down the group (H-F\ 562, H-Cl\ 431, H-Br\ 366, H-I\ 299\ mathrmkJ mol^-1) because the halogen atom becomes larger down the group, so the shared bonding pair lies further from the halogen nucleus and is held less strongly; a weaker bond requires less thermal energy to break, so HI decomposes most readily.
🧑🏫 Teacher's explanation
Think of the bond between Hydrogen and a Halogen like a magnet holding two balls together. If the balls are small and close together, the magnetic pull is incredibly strong and hard to break. As you move down Group 17, the halogen atom gets much larger. This means the 'center' of the electron cloud is much further away from the hydrogen nucleus. Because the distance increases, the electrostatic attraction between the nuclei and the shared pair of electrons weakens significantly.
📖 Definitions
🔍 Worked example — Easy
Calculate the decrease in bond enthalpy when moving from HCl to HBr, given H-Cl = 431 kJ mol^-1 and H-Br = 366 kJ mol^-1.
🔍 Worked example — Standard
A sample of HI(g) is heated. The reaction is 2HI(g) leftharpoons H_2(g) + I_2(g). If the bond enthalpy of H-I is 299 kJ mol^-1, how much energy is required to break 2 moles of H-I bonds?
🔍 Worked example — Exam-style
Explain why HF is much more thermally stable than HI using the concept of atomic radius and bond enthalpy.
⚠️ Common errors
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🎯 Exam question
Key ideas — Displacement reactions and comparing oxidising strength
Because oxidising strength falls down the group, a halogen higher in the group can oxidise (and so displace) the halide ion of one lower in the group from aqueous solution, but not the reverse. Chlorine water added to potassium bromide solution turns it orange-brown as bromine is formed: Cl_2(aq) + 2KBr(aq) arrow 2KCl(aq) + Br_2(aq). Chlorine water added to potassium iodide solution turns it brown (then almost black with excess) as iodine is formed: Cl_2(aq) + 2KI(aq) arrow 2KCl(aq) + I_2(aq). Bromine water added to potassium iodide solution similarly displaces iodine. No visible reaction occurs if bromine is added to potassium chloride solution, or iodine to potassium bromide or potassium chloride solution, since a weaker oxidising agent cannot remove electrons from a halide ion that is already a weaker reducing agent than itself would require.
🧑🏫 Teacher's explanation
Think of halogens as 'electron thieves'. In chemistry, an oxidising agent is something that wants to steal electrons from other substances. The higher up Group 17 you go (like Chlorine), the more 'hungry' the atom is for electrons because its nucleus has a very strong pull on them.
📖 Definitions
🔍 Worked example — Easy
Predict the products and state symbols for the reaction between chlorine gas and aqueous potassium bromide: Cl_2(g) + 2KBr(aq) arrow ?
🔍 Worked example — Standard
A student adds iodine solution to a solution of sodium chloride. State whether a reaction occurs and explain why.
🔍 Worked example — Exam-style
Explain why HCl is more thermally stable than HI using bond enthalpy arguments.
⚠️ Common errors
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🎯 Exam question
(a) State the observation made. [1] (b) Write a balanced chemical equation for the reaction. [2] (c) Explain why no reaction occurs if iodine is added to a solution of potassium bromide. [2] [5 marks]
Worked example
Question. Chlorine water is added to separate aqueous solutions of potassium bromide and potassium iodide.
🔍 Worked example — Easy
Predict the observation when aqueous bromine (Br_2) is added to a solution of potassium chloride (KCl).
🔍 Worked example — Standard
Write the balanced ionic equation for the reaction between chlorine (Cl_2) and potassium iodide (KI).
🔍 Worked example — Exam-style
Explain why Cl_2 can displace I^- from KI, but I_2 cannot displace Cl^- from KCl. Use the concept of electronegativity/oxidising strength.
⚠️ Common errors
Common mistakes
Practice
Answers