Why does chlorine oxidise bromide ions to bromine, while iodine cannot similarly displace chlorine from chloride ions? Group 17 halogens are oxidising agents: a halogen molecule can accept electrons and form halide ions. Their oxidising ability generally decreases down the group from chlorine to bromine to iodine under the standard school comparison. A JC2 H2 Chemistry student must connect displacement observations to redox half-equations, standard electrode potentials and the actual species present—not simply memorise that orange or brown means ‘a reaction happened’.
For JC2 Chemistry Tutor Hougang readers, the Singapore–Cambridge 2027 H2 Chemistry 9476 syllabus explicitly tests relative Group 17 oxidising power using standard reduction potentials, changes in physical state and volatility, and the thermal stability of hydrogen halides. The associated Qualitative Analysis Notes also give characteristic halogen colours in aqueous solution and in hexane, although the 2027 practical examination does not require students to carry out tests involving hexane. This guide separates examinable chemical reasoning from unnecessary practical assumptions.
- Cl₂ is the strongest oxidising agent of the three halogens in the ordinary comparison
- Cl₂ displaces Br⁻ and I⁻; Br₂ displaces I⁻
- Halide reducing power reverses the order: I⁻ > Br⁻ > Cl⁻
- Explain volatility through intermolecular London dispersion, not oxidising power
- Use actual E° values and supplied observation conditions
The precise 2027 H2 periodicity scope
The official SEAB H2 Chemistry 9476 syllabus for 2027 covers chlorine, bromine and iodine as Group 17 examples. It requires students to use standard E° values to deduce relative oxidising power, explain volatility in terms of instantaneous dipole–induced dipole attractions, and discuss hydrogen-halide thermal stability using bond energies. Students may also identify unknown elements from observed periodic properties.
These tasks overlap concepts in electrochemistry and practical observation, but they are not the same calculation. Oxidising power concerns electron acceptance, volatility concerns intermolecular attractions, and hydride thermal stability concerns covalent H–X bond strength. Keep those explanatory models separate before linking them.
A halogen molecule acts as an oxidising agent
A typical Group 17 halogen molecule X₂ can accept two electrons, giving two halide ions: X₂+2e⁻→2X⁻. This is reduction of the halogen. Because the halogen accepts electrons from another substance, it acts as the oxidising agent for that partner. In the corresponding displacement, a halide ion loses electrons and becomes its neutral halogen molecule.
A student who calls chloride ion the oxidising agent because its name contains chlorine has confused the elemental and ionic species. Cl₂ and Cl⁻ have different oxidation states and different electron roles. Write the charges and half-equations explicitly.
The reduction potential table gives a reliable order
Illustrative approximate standard potentials are E°(Cl₂/Cl⁻)=+1.36 V, E°(Br₂/Br⁻)=+1.07 V and E°(I₂/I⁻)=+0.54 V under the specified standard conditions. A more positive standard reduction potential indicates a stronger tendency for that reduction relative to the standard hydrogen electrode under the model. This supports the order Cl₂ > Br₂ > I₂ in oxidising power.
Students should use the actual Data Booklet values in an examination rather than memorise rounded figures. A standard potential comparison is a thermodynamic prediction under standard conditions, not a claim about every solution concentration or the immediate rate of a visible colour change.
Chlorine can oxidise bromide
The overall net ionic equation is Cl₂+2Br⁻→2Cl⁻+Br₂. Chlorine is reduced from oxidation number zero to −1, while bromide is oxidised from −1 to zero. This reaction is the classic displacement in which the stronger oxidising halogen produces the weaker oxidising halogen from its ions under suitable conditions.
To explain the displacement, a student should write which species receives the electrons and which loses them. The fact that chlorine lies above bromine in Group 17 is a convenient trend clue, but the oxidation and reduction half-equations give the chemically complete reasoning.
The half-equations reveal the two-electron balance
For the chlorine–bromide reaction, reduction is Cl₂+2e⁻→2Cl⁻ and oxidation is 2Br⁻→Br₂+2e⁻. Adding them cancels the electrons and gives the balanced net ionic equation. Both charge and element counts agree. The oxidising agent is chlorine; the reducing agent is bromide.
A student who writes Cl₂+Br⁻→Cl⁻+Br₂ has omitted a coefficient and fails both electron and atom conservation. The simple two-electron balance is worth checking before deciding on an observation.
Calculate the E°cell for chlorine displacing bromide
With illustrative standard reductions of +1.36 V for Cl₂/Cl⁻ and +1.07 V for Br₂/Br⁻, the chlorine reduction is the cathodic reduction and bromide is oxidised. E°cell=+1.36−(+1.07)=+0.29 V for the reaction as written. The positive sign is consistent with standard thermodynamic favourability.
Do not add 1.36+1.07 merely because one reaction is oxidation and the other reduction unless you have correctly changed the sign of the reversed half-reaction potential. Writing cathode reduction minus anode reduction is often the simplest way to preserve the standard table convention.
Chlorine can also oxidise iodide
The net ionic equation is Cl₂+2I⁻→2Cl⁻+I₂. Chlorine gains two electrons and iodide loses two. Under the approximate standard values, E°cell=1.36−0.54=+0.82 V. This is a larger positive difference than for the chlorine–bromide comparison under the same standard model.
The greater positive standard potential does not directly tell the exact speed or intensity of the colour change. It represents a thermodynamic difference between the specified redox couples. Actual observations depend on solution concentration and kinetics as well as the chemical products.
Bromine can oxidise iodide
The net ionic equation is Br₂+2I⁻→2Br⁻+I₂. The standard cell potential estimated from +1.07 and +0.54 V is about +0.53 V, supporting forward displacement under standard assumptions. Iodide is the reducing agent because it donates electrons as it becomes iodine.
This example is valuable because it shows that a halogen does not need to be the most powerful oxidising agent in the entire group to displace another halogen. It only needs sufficient oxidising tendency relative to the halide partner in the stated conditions.
Iodine does not displace bromide in the standard model
For the hypothetical reverse reaction I₂+2Br⁻→2I⁻+Br₂, the standard potential is approximately 0.54−1.07=−0.53 V. That direction is not thermodynamically favoured under the standard comparison. The reverse bromine-plus-iodide direction has the positive value.
A student who sees iodine’s dark colour and concludes it must be more reactive than bromine has used appearance instead of electrochemical evidence. The standard reduction potentials supply the relevant oxidising-power comparison.
The two-by-three displacement prediction table
With Cl₂, Br₂ and I₂ as possible reagents and Cl⁻, Br⁻ and I⁻ as possible dissolved halides, the ordinary predictions are: Cl₂ displaces Br⁻ and I⁻; Br₂ displaces I⁻; I₂ displaces neither Cl⁻ nor Br⁻. A halogen with its own matching halide does not create a new net displacement product in this simple classification.
Do not confuse the possibility of iodine–iodide complex equilibria with the redox displacement of another element. A solution can change apparent colour through speciation without necessarily involving a stronger halogen oxidising a weaker halide.
The halide reducing-power trend is opposite
Since I⁻ is most readily oxidised of the three ordinary halide ions in the comparison, it is a stronger reducing agent than Br⁻, which is stronger than Cl⁻ under the relevant standard redox ranking. A strong oxidising halogen is therefore paired with a conjugate halide ion that is relatively weak as a reducing agent, and vice versa.
Students sometimes use the Cl₂ > Br₂ > I₂ oxidising order to rank Cl⁻ > Br⁻ > I⁻ as reducing agents without reversing it. The species and electron directions differ, so the order reverses. The two half-reaction arrows provide a useful memory anchor.
Aqueous colours help but are not infallible
The 2027 Qualitative Analysis Notes describe chlorine’s aqueous solution as pale yellow, bromine’s as orange and iodine’s as brown in the stated comparison. These visual observations can support a reaction inference when concentration and context are suitable, but they do not substitute for balanced redox chemistry. Dilution, mixtures and chemical equilibria can alter apparent colour.
A student who writes ‘orange means bromide ions’ has reversed the product identity. Bromide ion solutions are ordinarily colourless in the simple school context; orange colour can be consistent with dissolved molecular bromine under the specified observations.
Iodine in water and in a nonpolar solvent
The syllabus’s halogen observation notes describe iodine as brown in aqueous solution and purple in hexane under the stated simple comparison. The difference arises from the molecular environment and visible absorption, and the aqueous colour can also be affected by iodide/triiodide equilibria. The colour is not a change in iodine’s elemental identity.
The important examination boundary is that the 2027 H2 practical syllabus says candidates will not be required to carry out tests involving hexane. Students can still interpret supplied theory data about hexane colours where provided; do not mistake reading a table for being required to perform a hexane extraction.
Bromine colour differs with medium
In the SEAB table, bromine is described as orange in aqueous solution and orange-red in hexane in the relevant comparison. Its appearance may vary with concentration and background conditions. The chemistry student should record whether the observation comes from a water layer, an organic layer or an unfamiliar prepared solution before naming a product.
If a theory problem supplies a two-layer observation, read the layers explicitly rather than writing a colour without identifying its location. But avoid inventing a hands-on extraction task that the practical syllabus excludes.
The physical states differ at room temperature
Chlorine is a greenish-yellow gas, bromine a reddish-brown liquid and iodine a dark solid that can produce purple vapour under appropriate conditions. These differences are part of the Group 17 physical trends in the ordinary school comparison. They do not mean the heavier elemental halogen automatically has a stronger oxidising effect.
Compare the actual chemical questions: state at room temperature concerns physical properties; displacement concerns reduction potential. A student who ranks oxidising power from boiling point alone is using the wrong evidence.
London dispersion attractions explain volatility
Cl₂, Br₂ and I₂ are nonpolar diatomic molecules but still experience instantaneous dipole–induced dipole attractions. Going down the group, larger and more polarisable electron clouds strengthen these dispersion interactions, so boiling temperatures tend to increase and volatility decreases. This accounts for the room-temperature progression gas, liquid, solid in the familiar sequence.
A student should not claim that iodine’s I–I covalent bond must be the strongest because iodine is solid. Melting or boiling molecular iodine primarily overcomes intermolecular attractions rather than breaking every I–I bond. Intramolecular bond strength and intermolecular dispersion are different quantities.
Thermal stability of hydrogen halides is another trend
The syllabus asks for thermal stability of Group 17 hydrides in terms of bond energies. The H–Cl bond is stronger than H–Br, which is stronger than H–I under the ordinary school trend; as bond strength falls down the group, the hydrides are generally less thermally stable to bond breaking. This is different from halogen oxidising strength and volatility.
A student who argues HI is the most thermally stable merely because iodine is the heaviest atom has ignored the bond-energy data. Use H–X bond strength, not boiling point or molecular size alone, to explain thermal stability.
Hydrogen-halide acid strength is not the same question
Hydrogen halides can act as acids in water, but aqueous acidity depends on dissociation and solvation as well as the H–X bond. The named 2027 Group 17 thermal-stability outcome asks specifically for bond energy reasoning. A student should not answer thermal stability solely by ranking the pH of solutions with unspecified concentrations.
Different properties often share the same chemical family without following the same explanatory rule. A high-quality answer identifies what the question actually measures before choosing a trend.
A note about chloride, bromide and iodide tests
Suitable precipitation tests can produce different silver halide appearances—chloride white, bromide cream and iodide yellow under appropriate conditions. This is not the same as observing chlorine, bromine or iodine molecules in a displacement experiment. The species are chemically different and the reagents serve different purposes.
Students should not confuse the orange colour of Br₂(aq) with the cream appearance of AgBr(s). One is an aqueous molecular halogen observation; the other is a precipitate formed from bromide ions and silver ions. State the test reagent and phase before identifying the substance.
Iodine and triiodide can complicate aqueous observations
In the presence of iodide, molecular iodine can form triiodide through I₂+I⁻⇌I₃⁻. This can change iodine’s apparent solubility and visible colour in aqueous systems. The existence of that equilibrium does not change the fundamental redox result that bromine can oxidise iodide to iodine in the standard displacement model.
A student should avoid inventing a uniquely defined I₂ concentration solely from colour intensity without relevant calibration or equilibrium information. Qualitative observations support deductions, but they do not automatically provide exact species amounts.
E°cell does not specify reaction speed
A positive standard cell potential supports thermodynamic favourability for the reaction direction as defined, but does not quantify how rapidly a visible halogen product appears. Concentration, mixing and kinetic barriers can influence observations. A larger E°cell also should not be converted directly into a percentage yield without a defined equilibrium model.
This is particularly important in data-based questions that compare standard potentials with practical observations. Distinguish the question ‘will this redox direction be favoured?’ from ‘how long until the solution turns orange?’
Temperature and concentration can change real conditions
Standard potentials are defined under particular reference conditions. The actual redox tendency in a mixture can depend on ion concentrations and environment, so a table-based standard prediction is not a universal statement about every possible nonstandard solution. The 2027 H2 syllabus requires qualitative understanding of concentration effects on electrode potentials, but not detailed Nernst-equation calculations.
A student should state the standard assumption when using E° values and not pretend to compute a precise nonstandard cell voltage without the relevant formula or supplied data. The right level of explanation is both correct and syllabus-aligned.
An unknown Group 17 element problem
Suppose an unknown diatomic halogen X₂ displaces iodide but not bromide under corresponding school conditions. If X is one of chlorine, bromine or iodine, bromine is a plausible candidate because it oxidises iodide but not bromide and is not strong enough to displace chloride. Additional physical properties such as room-temperature state or colour can support the conclusion.
The answer must use both positive and negative displacement evidence. A student who identifies chlorine from ‘can displace iodide’ alone ignores that chlorine would also displace bromide, contradicting the other observation.
A six-cell redox audit
Make a small matrix using Cl₂, Br₂ and I₂ against chloride, bromide and iodide. For each distinct pair, predict whether net displacement is favoured using the relative standard potential order. Then write the balanced ionic equations for the three favourable distinct combinations. The matrix should include negative predictions with reasons, not only the colourful positive tests.
After practising, replace element names with unknown labels and provide E° data. A student who can infer direction from the potentials rather than the familiar names has acquired a transferable A-Level method.
Safety and actual practical requirements
Halogens and related solutions can be hazardous; chlorine and bromine pose significant risks, while organic solvents can add flammability and exposure concerns. This article contains no instructions to mix reagents or perform halogen displacement at home. The official 2027 practical paper excludes required tests involving hexane, and any laboratory experiments belong to authorised school supervision.
A written tutorial can use supplied observation tables, balanced equations and E° values to teach the scientific reasoning without exposing students or families to hazardous materials. Parents should ask for an accurate account of practical support rather than assume private tutoring includes chemical laboratory access.
Case 1: chlorine versus bromide
The reaction Cl₂+2Br⁻→2Cl⁻+Br₂ transfers two electrons. Chlorine is reduced, bromide oxidised, and the positive illustrative E°cell≈0.29 V supports the forward displacement under standard conditions. A student should be able to establish the conclusion from half-equations and reduction potentials, not solely from an orange colour appearing in a fictional observation.
Change the halide from bromide to iodide and calculate the revised cell potential using the relevant E° values.
Case 2: chlorine versus iodide
With Cl₂+2I⁻→2Cl⁻+I₂, the standard potential difference is approximately 1.36−0.54=+0.82 V using the rounded values. The reaction produces elemental iodine from iodide ions, but its apparent colour can depend on solution composition and iodine species. The oxidation state change is still I from −1 to 0 and Cl from 0 to −1.
Write the two individual half-equations and check total ionic charge.
Case 3: bromine versus iodide
Br₂+2I⁻→2Br⁻+I₂ is the familiar positive displacement direction. From illustrative values +1.07 V and +0.54 V, E°cell is +0.53 V. Bromine accepts electrons as oxidising agent while iodide supplies them as reducing agent. Merely saying ‘iodine is lower in the table’ is less complete than describing the electron transfer.
Replace iodide with chloride and decide whether the same bromine-based forward displacement is favoured.
Case 4: iodine versus bromide
For I₂+2Br⁻→2I⁻+Br₂, the potential difference is 0.54−1.07=−0.53 V. The negative sign indicates this direction is not favoured under the chosen standard model. Students who draw the forward product and then try to force a positive sign have ignored the reduction-potential order. The reverse reaction is favoured under matching assumptions.
Reverse the equation explicitly and explain why the potential sign changes.
Case 5: iodine versus chloride
An iodine molecule is a weaker standard oxidising agent than chlorine. It cannot displace chlorine from chloride in the corresponding standard school prediction. Any observed colour change in an iodine-containing chloride solution would need another explanation, such as solvent effects or equilibrium speciation, rather than a fabricated Cl₂ product.
Given a brown solution containing iodine and chloride, identify why colour alone is insufficient to claim oxidation of chloride.
Case 6: bromide and bromine are different chemical species
Bromide Br⁻ has oxidation state −1 and is usually colourless in the ordinary aqueous ion context, while elemental Br₂ contains bromine at oxidation state zero and can impart an orange appearance under the stated conditions. A student who says orange indicates bromide has confused a dissolved ion with its oxidised elemental product. The species label matters.
Write the oxidation half-equation for bromide and identify which form supplies orange colour in the standard table.
Case 7: aqueous iodine versus iodide
Iodide I⁻ is a halide ion and does not automatically produce the familiar brown iodine appearance. Elemental iodine I₂ in aqueous solution can produce brown colour under the school comparison, and in the presence of iodide the I₂/I₃⁻ equilibrium may alter speciation. The colours therefore should be interpreted with chemical context and not treated as unique molecular fingerprints.
Compare a solution initially containing KI with one containing I₂, and state which iodine oxidation state each species has.
Case 8: the triiodide equilibrium is not redox displacement
I₂+I⁻⇌I₃⁻ forms triiodide by association of iodine with iodide. The equilibrium can affect apparent aqueous iodine solubility and colour, but it is not the same net reaction as Cl₂ oxidising I⁻ to I₂ with chloride production. Students should distinguish ligand-like association equilibria from a process with clear oxidation-number changes.
Ask which quantities conserve iodine atoms and net charge in the I₃⁻ equation, then compare with chlorine–iodide redox.
Case 9: a hexane colour is not a required Paper 4 manipulation
The 9476 qualitative notes include halogen colours in hexane, but the 2027 practical syllabus states candidates will not be required to carry out tests involving hexane. A written question may still provide data about nonpolar-solvent colours as information to interpret. Students should learn the chemical principles without mistaking a reference table for a required hazardous extraction experiment.
Explain why a theory observation in a supplied table can be examined without students actually performing that practical test.
Case 10: orange versus orange-red depends on medium
The notes describe bromine as orange in an aqueous solution and orange-red in a hexane layer. Differences in solvent environment affect observed absorption and colour. A learner must read which phase was observed before using the colour in a deduction. It is not valid to claim one exact shade occurs at every concentration or with every background substance.
In a data-based question, label each observed phase and distinguish the colours from those of dissolved halide ions.
Case 11: iodine’s purple vapour is not liquid bromine
Solid iodine can produce purple vapour under suitable heating, while bromine is a reddish-brown liquid near room temperature and chlorine is a greenish-yellow gas. These observations help identify physical states, but they are not measures of oxidising power. A student should not conclude purple vapour means iodine must have the highest E° value.
Rank halogen oxidising power using standard potentials separately from ranking room-temperature volatility.
Case 12: polarity and boiling point are separate
Cl₂, Br₂ and I₂ are nonpolar diatomic molecules but possess transient electron-density fluctuations, giving London dispersion forces. As the electron cloud becomes larger and more polarisable down the group, dispersion attractions grow and volatility generally falls. The molecules need not have a permanent electric dipole to attract one another.
Explain the gas-to-liquid-to-solid progression without claiming new ionic bonds form between halogen molecules.
Case 13: covalent bond strength does not explain boiling directly
Breaking an X–X bond requires energy, but boiling an elemental halogen generally separates molecules without dissociating every molecule into atoms. Therefore room-temperature physical state is chiefly explained by intermolecular attractions rather than the covalent bond dissociation energy alone. A claim that solid iodine has the strongest I–I bond because it is solid confuses processes.
Draw one I₂ molecule and two neighbouring I₂ molecules, then label which interaction is relevant to boiling.
Case 14: volatility and oxidising power trend in different directions
Chlorine is the most volatile and strongest standard oxidising agent of Cl₂, Br₂ and I₂ in the ordinary comparison, while iodine is least volatile and weakest oxidising agent. Those facts arise from different physical causes: intermolecular dispersion versus electron-accepting redox thermodynamics. They should not be conflated into one law of atomic size.
Given two trend questions, select the correct explanation for each rather than reusing one generic sentence.
Case 15: hydrogen chloride versus hydrogen iodide thermal stability
The H–Cl covalent bond is stronger than H–I in the ordinary Group 17 hydride comparison, so HCl is more thermally stable to bond cleavage than HI under suitable conditions. This concerns intramolecular bond energy, not the boiling point of the elemental halogen molecule or the colour of its aqueous solution. A student should name the bond whose strength changes.
Compare H–Br as the intermediate case and state the expected qualitative trend without inventing exact bond energy numbers.
Case 16: halide reducing power reverses halogen oxidising power
The elemental oxidising power order is Cl₂ > Br₂ > I₂, while the simple halide-ion reducing power order is I⁻ > Br⁻ > Cl⁻. The trend reversal follows the paired reduction and oxidation directions. A student who ranks chloride as the most powerful reducing agent by copying chlorine’s position has used the wrong chemical species.
Write Cl₂/Cl⁻ and I₂/I⁻ as paired redox couples and mark which member accepts and which donates electrons.
Case 17: E° difference determines direction, not equation coefficients
When Cl₂ accepts two electrons and Br⁻ supplies them, both half-equations have two electrons, but the potential difference is calculated as 1.36−1.07, not 2(1.36−1.07). Electrode potential is an intensive quantity and should not be multiplied by balanced electron coefficients. The two-electron count matters for ΔG°=−nFE°, which is a different quantity.
Ask the learner to compute the sign and magnitude of E°cell without scaling it with stoichiometric coefficients.
Case 18: the energy relationship can use n=2
For Cl₂+2Br⁻→2Cl⁻+Br₂, the idealised standard cell potential is about +0.29 V and two electrons transfer per reaction as balanced. The relation ΔG°=−nFE° would give an illustrative negative standard free-energy change for the forward reaction. This links Group 17 displacement to wider H2 electrochemistry, while keeping potentials and extensive energy distinct.
Use F≈96485 C mol⁻¹ and n=2 to estimate the standard energy change per mole of balanced reaction if requested.
Case 19: a positive standard potential is not a percentage yield
A displacement reaction with positive E°cell is thermodynamically favoured under its standard conditions, but that number does not alone give an exact final percentage conversion of a finite mixture. The initial amounts, nonstandard concentrations, equilibrium behaviour and kinetics can affect observations. Students should not turn +0.82 V into an invented 82% iodine yield.
Identify what extra data a quantitative equilibrium or limiting-reagent calculation would need.
Case 20: the negative prediction is useful evidence
If an unknown halogen displaces iodide but not bromide under suitable standard comparisons, it may be bromine within the Cl/Br/I candidate group. Chlorine would displace both, while iodine would not oxidise iodide in that sense to create a new product. The negative observation helps narrow the identity; one positive result alone is insufficient.
Design a two-test decision map that distinguishes Cl₂, Br₂ and I₂ by the given halide-displacement tendencies.
Case 21: a halide test is not a displacement test
Adding suitable silver ions to a halide solution can produce AgCl white, AgBr cream or AgI yellow precipitates under relevant conditions. That is precipitation chemistry involving ionic salts, not oxidation of a halide into its elemental halogen. A student should distinguish a solid AgBr observation from orange dissolved Br₂ formed after oxidation.
Given ‘cream precipitate’ and ‘orange solution’ as two observations, identify which test family each belongs to.
Case 22: temperature and concentrations matter to interpretation
E° values are defined at standard conditions and cannot be blindly treated as exact potentials for every solution concentration and temperature. The H2 syllabus expects qualitative awareness of concentration effects without requiring Nernst equations. In unfamiliar data-based problems, students should use supplied chemical conditions rather than claim a standard result is numerically identical to the observed nonstandard voltage.
State the standard assumption beside the potential difference before predicting direction.
Case 23: colour observation can be inconclusive
A pale yellow solution might contain a small amount of chlorine in the standard table, but other coloured species or mixtures can appear in practical settings. No one colour name establishes a unique chemical identity without supporting reagent and reaction evidence. A careful candidate reports the observed shade, phase and conditions, then links it to the appropriate balanced chemical possibility.
Ask what would make a proposed bromine identification more convincing than a vague orange observation alone.
Case 24: an integrated exam question
A new paper gives a halogen colour table, standard reduction potentials and two displacement outcomes from unknown labels X and Y. The student must first infer which halogen is the stronger oxidising agent, write charge-balanced half-equations, calculate an E°cell and explain whether an observed colour is consistent with the proposed product. Each result comes from distinct evidence and none may be assumed solely from the element name.
After feedback, replace chlorine, bromine and iodine with labels and changed E° values to test whether the reasoning transfers.
Case 25: a tutor should diagnose the first wrong idea
One student may calculate E° with the wrong subtraction direction, another may confuse bromide with bromine colour, while a third may explain iodine’s low volatility through a strong I–I covalent bond. Each learner needs a different correction: redox convention, species identification or intermolecular-force explanation. A correct group table copied from memory would not reveal those hidden misconceptions.
Set a new problem targeting each first error and check independent explanation a few days later.
Case 26: a safe final revision route
Session one covers halogen and halide identities, session two balanced displacement equations, session three E°cell calculations, and session four the opposite reducing-agent trend. Sessions five and six treat physical-state volatility and hydrogen-halide bond energies. The last session combines an unseen colour table with an explicit note about the lack of required hexane testing in Paper 4.
The goal is a student who can independently select the right chemical model for each question and avoid hazardous practical improvisation. No tutor can guarantee the final examination grade.
Frequently asked Group 17 questions
Which is the strongest oxidising halogen among Cl₂, Br₂ and I₂? Chlorine under the ordinary standard comparison. Can bromine oxidise iodide? Yes. Can iodine oxidise bromide under standard conditions? Not in the favoured direction. What is the reducing-power order for their halide ions? I⁻ > Br⁻ > Cl⁻. Why does iodine have a higher boiling point? Stronger dispersion attractions between larger, more polarisable molecules.
What colour is iodine in the SEAB aqueous comparison? Brown. Does orange automatically identify Br⁻? No; dissolved Br₂ is a different species. Does positive E°cell prove a fast reaction? No. Are hexane tests required hands-on in 2027 Paper 4? No. Why does HI have lower thermal stability than HCl in the bond-energy comparison? Its H–I bond is weaker.
Further reading and JC2 tutor enquiries
Continue with JC1 displacement foundations, Group 2 and Group 17 periodicity, standard electrode potentials and cell measurements, and the 2027 SEAB H2 9476 syllabus and Qualitative Analysis Notes for exact scope.
For current JC2 H2 Chemistry tutorials, contact eduKate Singapore or ask about current three-student group availability. Confirm tuition fees, timetable and travel to Fourth Avenue. The values and questions are original teaching examples.