The shape of the lesson
By the end of the lesson, learners can
- Displacement reactions: Add each halogen to each halide solution in the simulator.
- With hydrogen, and the H–X bond: Compare how the halogens react with hydrogen, then heat each hydrogen halide in the simulator.
- Oxidising power: Cl2 > Br2 > I2.
- Colours in hexane: Br2 orange-red, I2 purple.
- H–X bond energy falls, so HX stability falls.
- Halogen reduced; halide oxidised.
One picture every idea returns to
Halogens are oxidising agents: they gain electrons to form halide ions. Oxidising power decreases down the group (Cl2 > Br2 > I2) because larger atoms attract an incoming electron less strongly. So a halogen displaces any halogen below it from a solution of its halide. The halogens react with hydrogen less and less vigorously down the group, and the thermal stability of the hydrogen halides decreases from HCl to HI because the H–X bond gets longer and weaker.
Forty-five minutes, phase by phase
| Time | Phase | What happens in the room | Grouping |
|---|---|---|---|
| 0 to 5 min | Hook: predict | Open with the question in the title, Who takes the electron?. Chlorine pulls electrons from bromide and iodide ions; iodine cannot take them back. Put the figure on the board: it is the one picture every idea in the lesson returns to. | Whole class, then pairs |
| 5 to 17 min | Teach and check: Displacement reactions | Add each halogen to each halide solution in the simulator. Project tab 2 of the lab (Group 17 lab, displacement) and let a student drive it while the class predicts. | Whole class, then pairs |
| 17 to 30 min | Teach and check: With hydrogen, and the H–X bond | Compare how the halogens react with hydrogen, then heat each hydrogen halide in the simulator. Project tab 3 of the lab (Group 17 lab, hydrogen halides) and let a student drive it while the class predicts. | Think-Pair-Share |
| 30 to 36 min | Check | The four Stage 1 questions from the student lesson, answered on mini whiteboards or by show of fingers. Re-teach the one most get wrong. | Individual, then whole class |
| 36 to 43 min | Plenary | The exam-style questions from Stage 2, one at a time: thirty seconds alone, then a random call. The feedback lines in the key give the examiner's wording. | Individual, random call |
| 43 to 45 min | Exit and homework | Set the worksheet below (6 questions, 14 marks). Learners self-assess against the objectives before they leave. | Individual |
Protect the parts that carry the learning
The sequence carries about fifty minutes of material, so the lab is driven by one student while the class predicts rather than explored freely. Times are a guide.
Protected: the check at 30 minutes. It is the only individual read on whether the key idea landed, and it is not cut.
Cut first if the build overruns: the last check in With hydrogen, and the H–X bond, never the check or the plenary.
In a 60 minute block: run worksheet questions 1 to 3 in class as a Rally Coach set and model both the hardest worked example and its check live.
Think-Pair-Share
Pose the question, give thirty seconds of silent thinking, then a minute to compare with a partner before a few pairs share with the class. The silent minute is the part that cannot be cut.
The full structure, with its variations and the other twelve, is on the cooperative-learning page.
What to head off, and how
| Trap learners fall into | Teaching move that pre-empts it |
|---|---|
| Oxidising power decreases down the group, so iodine cannot oxidise chloride or bromide ions. | When explaining displacement, identify the halogen as the oxidising agent and the halide ion as the reducing agent. For thermal stability, refer to the H–X bond energy, not to intermolecular forces. |
Halogens as oxidising agents and the hydrogen halides: worksheet
6 questions, 14 marks. Written in the style of the papers; no past paper question is reproduced.
Write ionic equations for the reactions of (a) chlorine with iodide ions and (b) bromine with iodide ions.
Describe how you would show that chlorine is a stronger oxidising agent than bromine, giving observations.
Explain why the oxidising power of the halogens decreases from chlorine to iodine.
Describe and explain the trend in thermal stability of HCl, HBr and HI.
Compare the conditions needed for chlorine and for iodine to react with hydrogen.
Use oxidation numbers to show that Br2 + 2I− → 2Br− + I2 is a redox reaction.
Answer key and marking notes
- (a) Cl2 + 2I− → 2Cl− + I2.
- (b) Br2 + 2I− → 2Br− + I2.
- Add chlorine water to aqueous potassium bromide; it turns orange as Br2 forms.
- Add bromine water to aqueous potassium chloride; no reaction (stays the orange of the added bromine).
- Shaking with hexane confirms the halogen present (orange-red for Br2).
- Atomic radius and shielding increase down the group.
- The nucleus attracts an incoming electron less strongly, so the halogen gains electrons less readily.
- Thermal stability decreases from HCl to HI.
- The H–X bond energy falls (431, 366, 299 kJ mol−1) because the halogen atom is larger and the bond longer.
- HI decomposes to H2 and purple I2 with a hot glass rod; HCl does not.
- Chlorine reacts explosively with hydrogen in sunlight (or a flame).
- Iodine reacts only slowly when heated, and the reaction is reversible.
- Br goes from 0 to −1: reduced.
- I goes from −1 to 0: oxidised.
8 multiple-choice questions on this lesson
From the topic question bank. Assign them in the app for a marked score in your gradebook, or run the topic as a class game.
Which statement about the oxidising power of the halogens is correct?
- Aoxidising power decreases down the group (Cl2 > Br2 > I2), because the larger atoms attract an extra electron less strongly
- Boxidising power increases down the group (I2 > Br2 > Cl2), because the larger atoms have more electrons
- Coxidising power increases down the group, because the X-X bond gets weaker
- Dall the halogens have the same oxidising power, because each gains one electron to form X−
Chlorine water is added to colourless aqueous potassium bromide. What is seen, and what is the ionic equation?
- Athe solution turns orange (yellow when dilute); Cl2 + 2Br− → 2Cl− + Br2
- Bthe solution turns brown; Cl2 + 2Br− → 2Cl− + Br2
- Cno change, because bromine is a stronger oxidising agent than chlorine
- Dthe solution turns orange; Br2 + 2Cl− → 2Br− + Cl2
Which mixture of aqueous solutions shows NO reaction?
- Abromine water + potassium chloride
- Bchlorine water + potassium iodide
- Cbromine water + potassium iodide
- Dchlorine water + potassium bromide
Which description of the reactions of hydrogen with the halogens is correct?
- Afluorine reacts explosively even in the dark, chlorine explodes in sunlight, bromine needs heating, and iodine reacts slowly and reversibly on heating
- Biodine reacts explosively with hydrogen, and fluorine reacts only slowly on heating
- Call the halogens react explosively with hydrogen at room temperature in the dark
- Dchlorine and hydrogen react only when heated with a catalyst, but bromine explodes in sunlight
A red-hot glass rod is plunged into separate gas jars of HCl, HBr and HI. Only one gas gives dense purple fumes. Which gas is it, and why?
- AHI, because the H-I bond (299 kJ mol−1) is the weakest, so HI decomposes most easily into hydrogen and iodine vapour
- BHCl, because the H-Cl bond is the strongest, so it releases the most energy as it breaks
- CHI, because it is the most thermally stable hydrogen halide
- DHBr, because bromine vapour is purple
Chlorine water is added to aqueous potassium iodide, then a little cyclohexane is added and the tube is shaken. The layers separate. What colour is the upper (cyclohexane) layer?
- Apurple, because the iodine formed dissolves in the non-polar solvent
- Bbrown, because iodine is always brown in solution
- Corange, because bromine is formed
- Dpale green, because unreacted chlorine collects in the organic layer
For the reaction Br2(aq) + 2I−(aq) → 2Br−(aq) + I2(aq), which statement is correct?
- Abromine is reduced from 0 to −1 and iodide is oxidised from −1 to 0, so bromine is the oxidising agent
- Bbromine is oxidised from 0 to −1, so bromine is the reducing agent
- Ciodide is reduced from −1 to 0, so iodide is the oxidising agent
- Dno oxidation numbers change, so this is not a redox reaction
120 cm3 of chlorine, measured at room conditions, is bubbled into excess aqueous potassium iodide. What mass of iodine is formed? Cl2 + 2KI → 2KCl + I2 (Ar I = 126.9; molar gas volume = 24.0 dm3 mol−1)
- A1.27 g
- B2.54 g
- C0.635 g
- D30.5 g
Quiz key
A halogen acts as an oxidising agent by gaining electrons: X2 + 2e− → 2X−. Down the group the atoms are larger and the incoming electron is further from the nucleus and more shielded, so it is attracted less strongly; oxidising power therefore decreases, Cl2 > Br2 > I2. Saying it increases down the group is the classic error, contradicted by the displacement reactions (chlorine displaces iodine, never the reverse). More electrons do not make an atom attract another electron more strongly. Although the I-I bond is weaker, the decreasing attraction for electrons dominates. All form X−, but how readily they do so differs, so the powers are not equal.
Chlorine is a stronger oxidising agent than bromine, so it oxidises bromide ions to bromine: Cl2 + 2Br− → 2Cl− + Br2. Aqueous bromine is orange, or yellow when dilute. A brown solution is typical of aqueous iodine, formed when chlorine or bromine reacts with iodide. Bromine is the weaker oxidising agent, so the reaction does occur. The last equation is the reverse reaction, bromine oxidising chloride, which does not happen; the colour change is real but it is caused by the forward reaction.
Bromine is a weaker oxidising agent than chlorine, so it cannot oxidise chloride ions: bromine water + KCl stays orange with no reaction. Chlorine oxidises iodide (Cl2 + 2I− → 2Cl− + I2, brown solution). Bromine oxidises iodide (Br2 + 2I− → 2Br− + I2, brown solution). Chlorine oxidises bromide to orange bromine. Each of these three has a halogen higher in the group than the halide, so each reacts.
Reactivity with hydrogen decreases down the group, matching the fall in oxidising power: H2 + F2 explodes even in the cold and dark; H2 + Cl2 explodes in sunlight (UV light); H2 + Br2 needs heating (often with a platinum catalyst); and H2 + I2 ⇌ 2HI is slow and reversible even on heating. Reversing the order treats iodine as the most reactive, the classic error of thinking oxidising power increases down the group. Only fluorine reacts explosively in the dark. Chlorine reacts more readily than bromine, not less.
Thermal stability of the hydrogen halides decreases down the group because the H-X bond gets weaker (431, 366, 299 kJ mol−1), as the larger X atom overlaps less effectively with H. HI therefore decomposes readily on the hot rod: 2HI → H2 + I2, giving purple iodine vapour. HCl has the strongest bond, so it is the most stable and is not decomposed; breaking a bond takes in energy, it does not release it. Calling HI the most stable reverses the trend, the classic error. Bromine vapour is orange-brown, not purple, and HBr decomposes only slightly at this temperature.
Chlorine oxidises iodide: Cl2 + 2I− → 2Cl− + I2. Iodine is non-polar, so on shaking it moves into the cyclohexane layer, where it is purple (the same colour as iodine vapour). In water, or in aqueous KI, iodine appears brown, which is why the organic layer is used to confirm iodine. There is no bromide present, so bromine (orange in cyclohexane) cannot form. Chlorine is the oxidising agent and is used up by the excess iodide; the purple colour shows iodine, not chlorine.
Bromine goes from 0 in Br2 to −1 in Br−: a decrease in oxidation number, so bromine gains electrons and is reduced; it is the oxidising agent. Iodine goes from −1 in I− to 0 in I2: an increase, so iodide is oxidised and is the reducing agent. A change from 0 to −1 is reduction, not oxidation. Going from −1 to 0 is oxidation, not reduction. Oxidation numbers clearly change, so the reaction is redox.
n(Cl2) = 0.120 / 24.0 = 0.00500 mol, and the ratio Cl2 : I2 is 1 : 1, so n(I2) = 0.00500 mol. Mr(I2) = 253.8, so the mass is 0.00500 × 253.8 = 1.27 g. 2.54 g doubles the moles, using the 2 in front of KI as if it applied to I2. 0.635 g uses the Ar of I (126.9) instead of the Mr of I2. 30.5 g multiplies the volume in dm3 by the Mr without dividing by the molar gas volume.
What to project
The student lesson embeds these tabs. Open them full screen on the projector; every slider has a drag handle on the stage.
Slides for this lesson
A projectable deck for every pack is on the way. Until then, present the student lesson full screen: its sections, figure and lab tabs follow the same order as the plan above.