Did you know? Chlorine can oxidise iodide to iodine, but the reversed iodine-plus-chloride reaction is not favoured under standard conditions. The SEAB data booklet explains why numerically: E°(Cl₂/Cl⁻)=+1.36 V versus E°(I₂/I⁻)=+0.54 V.
This original Hougang JC1 H2 Chemistry guide follows the 2027 syllabus 9476 Group 17 oxidising-strength outcome. It develops oxidation numbers, half-equations, E°cell sign audits, pairwise chlorine/bromine/iodine predictions and the distinction between thermodynamic direction and reaction rate. All examples are paper chemistry rather than halogen-handling instructions.
The immutable eduKateSG Clementi small-group reference is retained unchanged. Worked voltages are drawn from the official SEAB Chemistry Data Booklet or labelled fictional transfer numbers.
EDUKATESG · HOUGANG · JC1 H2 CHEMISTRY
14 chapters · Electrode potential matrix · Transfer test · FAQs · SEAB sources
Full chapter index: redox before colour
- 1. The asymmetric result is about redox driving force, not just colour
- 2. Identify which species donates and accepts electrons
- 3. Write the reduction and oxidation half-equations separately
- 4. Read the SEAB Group 17 standard reduction potentials
- 5. Calculate the standard cell potential for chlorine oxidising iodide
- 6. Show why iodine cannot displace chloride in the reverse standard direction
- 7. Use the bromine intermediary to construct the complete ranking
- 8. Explain why halide reducing strength increases in the opposite direction
- 9. Interpret colour evidence without replacing redox chemistry
- 10. Keep standard conditions separate from arbitrary solution concentrations
- 11. Distinguish halogen displacement from H–X thermal stability
- 12. Optional extension: relate positive E° to free energy
- 13. Solve unfamiliar displacement predictions without a memorised list
1. The asymmetric result is about redox driving force, not just colour
Chlorine is a stronger oxidising agent than iodine in the ordinary Group 17 comparison. Consequently, chlorine can oxidise iodide ions into iodine while itself being reduced to chloride. The balanced net ionic transformation is Cl₂ + 2I⁻ → 2Cl⁻ + I₂, with appropriate physical states assigned for the question.
Now reverse the chemical roles: iodine would have to oxidise chloride ions to elemental chlorine while being reduced to iodide. Under standard aqueous redox conditions this is not thermodynamically favoured. The failure of the reverse displacement is not a consequence of an incorrectly chosen balancing coefficient; it is a different direction of electron transfer with an unfavourable standard cell potential.
The 2027 Singapore H2 Chemistry 9476 Periodic Table section specifically requires using standard electrode potentials to describe and deduce the relative oxidising reactivity of Group 17 elements from chlorine to iodine. This guide develops that evidence using the SEAB Chemistry Data Booklet rather than memorising a reaction list without the underlying half-equations.
Any colour observation is secondary to the chemistry. Iodine can give a characteristic brown aqueous appearance or a purple colour in a nonpolar medium, but a solution’s exact shade depends on composition and conditions. The crucial evidence is the redox equation and the sign of the calculated potential.
Student question: Why can chlorine liberate iodine from iodide in the school redox model?
Answer standard: The Cl₂/Cl⁻ couple has a higher standard reduction potential than I₂/I⁻, giving a positive overall E°.
Common trap: Do not say chlorine simply replaces iodine because its atomic number is smaller.
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2. Identify which species donates and accepts electrons
An oxidising agent accepts electrons and is reduced. A reducing agent donates electrons and is oxidised. In the Cl₂ plus iodide reaction, each chlorine atom goes from oxidation state zero in Cl₂ to −1 in Cl⁻. The iodine in I⁻ goes from −1 to zero in I₂.
The donor is iodide, so iodide is the reducing agent. The acceptor is chlorine, so chlorine is the oxidising agent. These roles are consistent with the net ionic reaction: two electrons released by two I⁻ ions are accepted by one Cl₂ molecule. No electrons appear in the final balanced full equation because they cancel when half-reactions are added.
Students often call chlorine the reducing agent because it ’causes’ iodine to appear, or call iodide the oxidising agent because iodine ends up as an element. This reverses the definitions. The proper question is which species gains or loses electrons, not which product has the more visually striking colour.
A tutor can ask the learner to record oxidation numbers in an original two-column ledger before writing ‘oxidising’ or ‘reducing’. Once the electron change is explicit, the chemical-role labels follow automatically.
Student question: Which reagent is oxidised in Cl₂ + 2I⁻ → 2Cl⁻ + I₂?
Answer standard: Iodide, whose iodine oxidation number rises from −1 to zero.
Common trap: The oxidising agent is reduced, even though its name contains the word oxidising.
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3. Write the reduction and oxidation half-equations separately
The chlorine reduction half-equation is Cl₂ + 2e⁻ → 2Cl⁻. It gains two electrons, one for each chlorine atom, and the electrical charge balances: −2 on both sides. The reverse iodine oxidation half-equation is 2I⁻ → I₂ + 2e⁻, with charge −2 on both sides as well.
Adding the half-equations cancels the two electrons. The combined equation is Cl₂ + 2I⁻ → 2Cl⁻ + I₂. The balancing can be audited by counting chlorine and iodine atoms and total charge on either side. It is a much more transferable method than copying a final equation from memory.
The standard electrode potential table lists reductions in a particular direction. When a half-equation is reversed to represent oxidation, do not automatically copy the tabulated positive sign into the corresponding oxidation potential without changing the convention appropriately. The convenient cell formula uses both tabulated potentials as reduction values and takes their difference.
A student should demonstrate this on a second halogen pair before concluding they understand the method. If the half-equations are not both balanced, a plausible-looking voltage calculation alone cannot repair the chemical equation.
Student question: What two half-equations combine to form chlorine–iodide displacement?
Answer standard: Cl₂+2e⁻→2Cl⁻ and 2I⁻→I₂+2e⁻.
Common trap: Do not add redox half-equations while leaving uncancelled electrons in the final net reaction.
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4. Read the SEAB Group 17 standard reduction potentials
The SEAB Chemistry Data Booklet reports standard reduction potentials for familiar halogen couples. In the usual reference convention, E°(Cl₂/Cl⁻)=+1.36 V, E°(Br₂/Br⁻)=+1.07 V and E°(I₂/I⁻)=+0.54 V. These values are numerical evidence for the halogen oxidising-strength sequence.
A larger, more positive standard reduction potential means a greater thermodynamic tendency for the oxidised member of that couple to accept electrons relative to the common standard hydrogen electrode reference. Therefore the standard oxidising-agent ranking for the named elements is chlorine above bromine above iodine.
The values describe the specified standard-state redox couples, not arbitrary measured rates or universal reaction speeds. Their actual chemical response in a mixture may depend on concentration, activities, solvent and kinetics. A positive standard potential provides thermodynamic direction under standard conditions but does not tell us an exact time for a colour change.
Students should name the couple and the direction of the listed half-reaction when using a data booklet. A raw voltage with no identified oxidised/reduced species is not a complete chemical justification.
Student question: Which halogen couple has the highest of the three E° values?
Answer standard: Cl₂/Cl⁻ at +1.36 V.
Common trap: Standard reduction potentials are not electronegativity values or bond dissociation energies.
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5. Calculate the standard cell potential for chlorine oxidising iodide
The calculation uses E°cell = E°(cathode reduction) − E°(anode reduction, as tabulated). Here chlorine is reduced, so its reduction potential is +1.36 V. Iodide is oxidised, so use the reduction potential of the reverse iodine couple, +0.54 V, as the quantity to subtract.
The resulting value is E°cell = 1.36 − 0.54 = +0.82 V. Because it is positive, the redox direction written as Cl₂ + 2I⁻ → 2Cl⁻ + I₂ is thermodynamically favourable under the associated standard-state assumptions.
Do not multiply either electrode potential by two simply because two electrons occur in both half-equations. Electrode potential is an intensive quantity; the electron count is handled in balancing and may enter a separate Gibbs-energy expression, but multiplying E° by stoichiometric coefficients would change the predicted voltage incorrectly.
The full reasoning can be written compactly in an examination: chlorine is the stronger oxidising agent, its E° reduction is more positive, and the overall cell potential is +0.82 V. Those statements mutually support one another.
Student question: What is E°cell for Cl₂ oxidising I⁻ from the SEAB values?
Answer standard: +0.82 V.
Common trap: Do not multiply an electrode potential by the number of electrons.
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6. Show why iodine cannot displace chloride in the reverse standard direction
Suppose a student writes I₂ + 2Cl⁻ → 2I⁻ + Cl₂ and expects it to proceed in the ordinary standard comparison. Now iodine is the species being reduced, and chloride is the species being oxidised. Subtracting the same reduction potentials in the new direction gives E°cell = 0.54 − 1.36 = −0.82 V.
A negative standard cell potential indicates that the reverse reaction is thermodynamically unfavourable under the defined standard conditions. That is why iodine cannot normally liberate chlorine from chloride solution simply by being mixed with it. This does not mean the equation is unbalanced—it is balanced but points against the favoured redox direction.
Many exam errors arise from checking only whether atoms and charge balance. Conservation is essential but insufficient to predict spontaneity. An atom-balanced equation may still describe an energetically disfavoured direction, so the standard electrode potential provides an independent physical check.
The word cannot should be interpreted within the specified chemical conditions and ordinary displacement comparison. External electrical energy or a different coupled process can drive reactions that would not be spontaneous as isolated standard cells. The syllabus question is about relative oxidising power, not a universal claim that chloride can never be converted to chlorine.
Student question: Why is I₂ + 2Cl⁻ → 2I⁻ + Cl₂ not the favoured standard displacement?
Answer standard: Its E°cell is −0.82 V, the reverse of the positive chlorine–iodide reaction.
Common trap: A balanced equation is not automatically thermodynamically favourable.
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7. Use the bromine intermediary to construct the complete ranking
Between chlorine and iodine sits bromine. The data show +1.36 V for chlorine, +1.07 V for bromine and +0.54 V for iodine. Consequently chlorine can oxidise bromide under standard conditions, and bromine can oxidise iodide. The reverse comparisons are unfavourable in the isolated standard redox model.
For Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂, E°cell = 1.36 − 1.07 = +0.29 V. For Br₂ + 2I⁻ → 2Br⁻ + I₂, E°cell = 1.07 − 0.54 = +0.53 V. Both are positive, yet the magnitudes differ. This provides a numerical hierarchy rather than only a colour-based mnemonic.
Students should not rank the elements alphabetically or simply assume a larger atom must have a larger standard reduction potential. Down this named series, changes in electron affinities, bond dissociation, hydration and other thermodynamic contributions jointly determine the actual values.
Try a three-row table giving each couple’s potential and asking learners to predict every possible pairwise displacement. A consistent student will obtain a positive potential only when the more powerful halogen oxidises the halide of the less powerful halogen.
Student question: Can Br₂ oxidise I⁻ under the standard model?
Answer standard: Yes; E°cell = 1.07−0.54=+0.53 V.
Common trap: The halfway halogen can oxidise iodide but does not ordinarily oxidise chloride.
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8. Explain why halide reducing strength increases in the opposite direction
If the neutral halogen X₂ is a strong oxidising agent, its reduced halide X⁻ tends to be a weaker reducing agent in the standard Group 17 comparison. Chlorine is the strongest of the three oxidisers, while iodide is the strongest of the corresponding halide reducing agents. Thus the reducing-strength order for the anions is I⁻ above Br⁻ above Cl⁻.
The two rankings are not contradictory. They refer to opposite electron-transfer directions. Chlorine’s tendency to accept electrons is strong, whereas chloride’s tendency to donate electrons to form chlorine is comparatively weak. Iodide, by contrast, is more readily oxidised to iodine by a suitable oxidant.
A student should write each half-equation and reverse it before discussing the opposite role. It is common to see the correct halogen ranking copied onto the halides as though both are ranked by the same property. That loses the distinction between reduction and oxidation.
The redox couple is the link between the element and its anion, but its oxidation and reduction members play different roles. Teaching both rankings together makes the relative-reactivity trend more meaningful.
Student question: Which of I⁻, Br⁻ and Cl⁻ is the strongest reducing agent in the named comparison?
Answer standard: I⁻.
Common trap: Strong oxidising power of Cl₂ does not imply Cl⁻ is a strong reducing agent.
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9. Interpret colour evidence without replacing redox chemistry
A familiar hypothetical qualitative observation uses an iodine-containing product as evidence of iodide oxidation. Iodine can appear brown in water and purple in a suitable nonpolar phase, while bromine and chlorine have their own characteristic colours. The SEAB qualitative-analysis notes give reference descriptions, but the exact appearance in a mixture depends on concentrations and other dissolved species.
Colour change alone cannot prove one unique equation when several coloured reactants or products are possible. For example, an initially pale halogen solution and an iodide-containing sample may become visually darker, but the redox assignment should still be supported by oxidation-number changes, appropriate half-equations and the relevant E° data.
This guide does not supply experimental mixing directions for toxic or corrosive halogens. The educational problem is to interpret observations provided on paper, not to reproduce halogen displacement experiments without professional laboratory safeguards.
An exam answer should describe a visible result only if the relevant observation or standard test is supplied or clearly specified. It should not invent a measured precipitate, colour timing or extraction yield when none was recorded.
Student question: Does a brown colour alone uniquely prove the chlorine–iodide equation occurred?
Answer standard: No. Chemical and redox context is required to support the species assignment.
Common trap: Do not confuse a qualitative indicator with a direct measurement of E°cell.
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10. Keep standard conditions separate from arbitrary solution concentrations
Standard electrode potentials are tabulated for specified reference conditions and activities. When concentrations or other conditions differ, the actual potential and the resulting direction of an equilibrium comparison can require more careful electrochemical analysis. The number +0.82 V is not a promise that every vessel at every concentration measures exactly that voltage.
The standard result still provides a reliable baseline for the familiar Group 17 oxidising-power sequence. The H2 Chemistry syllabus explicitly asks students to deduce relative reactivity using E° values, so the correct answer starts from those data rather than adopting an unsupported environmental generalisation.
A very strong shift in chemical composition can affect reaction quotient and effective driving force. This is a reason to state the standard assumptions, not an invitation to claim that any weak oxidant must become strong if arbitrary amounts of electrolyte are added.
Students should resist trying to infer kinetics from the standard potential magnitude alone. Positive E°cell describes a favourable thermodynamic direction under standard conditions, while actual rate can be governed by activation barriers, phases and transport.
Student question: Does +0.82 V specify how many seconds the reaction takes?
Answer standard: No; it is a standard thermodynamic redox potential difference.
Common trap: Thermodynamic favourability and reaction speed are separate questions.
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11. Distinguish halogen displacement from H–X thermal stability
Another Group 17 syllabus outcome concerns the thermal stabilities of HCl, HBr and HI, explained by decreasing H–X covalent bond strengths down the group. That is not the same property as oxidising power of Cl₂, Br₂ and I₂. The reactant types and relevant thermodynamic quantities differ.
For hydrogen halide stability, a tutor considers H–X bond dissociation energies and the effect of increasing halogen size on bond length and overlap. For halogen displacement, the tutor considers reduction potentials and electron transfer between neutral halogen and aqueous halide. Both topics appear within Group 17, but they answer different chemical questions.
An answer such as ‘chlorine displaces iodine because HCl has a stronger bond than HI’ skips the actual redox half-equations. H–X bond energy does not directly appear in the simple Cl₂/Cl⁻ and I₂/I⁻ electrode comparison in that form.
The student should be able to look at a question and decide whether it asks for oxidising power, halide reducing power, hydride stability or halogen volatility. Each has its own relevant evidence and cannot be solved by copying one remembered down-group trend.
Student question: Are HCl/HBr/HI bond strengths sufficient by themselves to calculate Cl₂/I⁻ E°cell?
Answer standard: No. Use the relevant standard halogen/halide electrode potentials.
Common trap: Do not collapse several Group 17 trends into one universal ‘reactivity’ ranking.
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12. Optional extension: relate positive E° to free energy
A more advanced thermodynamic link is ΔG°=−nFE°cell, where n is the number of transferred moles of electrons per stoichiometric reaction and F is the Faraday constant. For the chlorine–iodide equation n=2 and the standard potential difference is +0.82 V.
Using F≈96485 C mol⁻¹, the illustrative calculation gives ΔG°≈−2×96485×0.82 J mol⁻¹≈−1.58×10⁵ J mol⁻¹, or about −158 kJ per mole of reaction as written. The negative sign is consistent with the favoured standard redox direction. The value depends on the rounded potentials and conventions.
This is supplementary electrochemical depth, not a claim that the 2027 H2 Periodic Table learning outcome explicitly demands a Gibbs-energy conversion from E°. For its core objective, identifying the relevant couples, E° difference and sign is sufficient.
The Faraday constant is not a factor by which electrode potentials themselves should be multiplied to obtain a new potential. It belongs in the separate free-energy relationship, which has energy units and depends on transferred electron count.
Student question: What sign of ΔG° corresponds to a positive E°cell in this extension?
Answer standard: Negative ΔG°, indicating standard thermodynamic favourability.
Common trap: E° is not multiplied by two when combining half-cell potentials.
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13. Solve unfamiliar displacement predictions without a memorised list
An exam-style paper table gives E°(X₂/X⁻)=+1.20 V and E°(Y₂/Y⁻)=+0.75 V for two fictional halogens. The learner should infer that X₂ is the stronger oxidising agent and can oxidise Y⁻, with E°cell=1.20−0.75=+0.45 V under standard conditions. The reverse Y₂+X⁻ reaction would have −0.45 V.
The exercise can be completed without knowing the colours of the fictional elements. Balanced half-equations and redox potentials provide the logic. This tests whether a student has learned the method rather than merely memorised chlorine, bromine and iodine positions.
A second transfer question swaps the labels on a real halogen pair and asks the learner to explain the resulting negative potential. Another supplies a standard potential with changed concentration data and asks whether the standard prediction alone can quantify the real current. The correct answer states the appropriate missing information.
Parents can ask one useful question after a tutorial: ‘Which species accepts electrons, and which E° value is more positive?’ A student who can identify both and form the appropriate subtraction has a transferable chemical decision method.
Student question: If E°(X₂/X⁻)=+1.20 and E°(Y₂/Y⁻)=+0.75 V, can X₂ oxidise Y⁻ under standard conditions?
Answer standard: Yes; E°cell=+0.45 V.
Common trap: A more positive reduction potential refers to greater oxidising tendency of the oxidised species.
Worked evidence: Group 17 oxidising strength by E°cell
The standard reduction potentials use the SEAB Chemistry Data Booklet values. Overall redox directions are determined by subtracting the reduction potential of the couple being oxidised from that of the couple being reduced, without multiplying voltages by coefficients.
| Couple or reaction | E° input | Calculated/qualitative result | Interpretation |
|---|---|---|---|
| Cl₂/Cl⁻ | +1.36 V | Strongest listed halogen oxidiser | Reference for Cl₂ displacement |
| Br₂/Br⁻ | +1.07 V | Intermediate oxidiser | Can oxidise I⁻ but not Cl⁻ in ordinary standard model |
| I₂/I⁻ | +0.54 V | Weakest listed halogen oxidiser | Iodide strongest reducing halide of three |
| Cl₂+2I⁻→2Cl⁻+I₂ | 1.36−0.54 | E°cell=+0.82 V | Favoured standard direction |
| I₂+2Cl⁻→2I⁻+Cl₂ | 0.54−1.36 | E°cell=−0.82 V | Unfavoured reverse |
| Cl₂+2Br⁻→2Cl⁻+Br₂ | 1.36−1.07 | E°cell=+0.29 V | Favoured |
| Br₂+2I⁻→2Br⁻+I₂ | 1.07−0.54 | E°cell=+0.53 V | Favoured |
| Br₂+2Cl⁻→2Br⁻+Cl₂ | 1.07−1.36 | E°cell=−0.29 V | Unfavoured reverse |
| I₂+2Br⁻→2I⁻+Br₂ | 0.54−1.07 | E°cell=−0.53 V | Unfavoured reverse |
The oxidation-number test, half-equation balancing, potential subtraction and final sign must agree. A correct numerical sign without a correct redox role is not dependable understanding. A positive standard E° gives a favourable redox direction under standard conditions, not an exact observed colour or reaction time.
Four independent redox diagnostics
- For Cl₂+2I⁻→2Cl⁻+I₂, write half-equations and find +0.82 V from the official data.
- For Br₂+2I⁻→2Br⁻+I₂, obtain +0.53 V and identify I⁻ as the reducing agent.
- For I₂+2Cl⁻→2I⁻+Cl₂, obtain −0.82 V and explain why the reverse displacement is not the favourable standard direction.
- For fictional E°(X₂/X⁻)=+1.20 V and E°(Y₂/Y⁻)=+0.75 V, infer X₂ can oxidise Y⁻ with +0.45 V and write a balanced equation.
Tutor preparation checklist
- Confirm the 2027 H2 Chemistry 9476 Section 5 Group 17 E° reactivity outcome.
- Read Cl₂/Cl⁻, Br₂/Br⁻ and I₂/I⁻ reduction potentials accurately from the SEAB data booklet.
- Identify oxidising and reducing agents from electron transfer and oxidation numbers.
- Write and balance both half-equations before calculating E°cell.
- Subtract tabulated reduction potentials using a fixed cathode-minus-anode convention.
- Do not multiply electrode potentials by electron stoichiometric coefficients.
- Distinguish standard thermodynamic favourability from an observed reaction rate.
- Predict chlorine–bromide, bromine–iodide and reverse reactions without guessing from colour.
- Keep hydride thermal stability and halogen volatility separate from oxidising power.
- Finish with unfamiliar numerical redox couples and ask the learner to justify the resulting sign.
Frequently asked questions
Why can chlorine displace iodide?
Chlorine has a more positive Cl₂/Cl⁻ reduction potential than iodine’s I₂/I⁻ couple, so it can oxidise I⁻ while being reduced.
Why does iodine not ordinarily displace chloride?
The reversed standard redox equation has E°cell=0.54−1.36=−0.82 V and is thermodynamically unfavourable as an isolated standard reaction.
What are the SEAB standard potentials for Cl₂, Br₂ and I₂ couples?
+1.36 V, +1.07 V and +0.54 V respectively, in their standard reduction directions.
Which Group 17 halogen is the strongest oxidising agent of the three?
Chlorine, followed by bromine and iodine.
Which halide ion is the strongest reducing agent?
Iodide, followed by bromide and chloride in the named comparison.
Why is E°cell not multiplied by electron coefficients?
Electrode potential is intensive; balance electrons in the half-equations and subtract the tabulated reduction potentials.
Does positive E°cell guarantee an immediate visible reaction?
No. It supports a thermodynamically favoured standard direction, but actual rates and observations depend on kinetic and solution conditions.
Is chlorine disproportionation in cold and hot alkali explicitly required in 2027 H2 9476?
The named Periodic Table Group 17 outcome instead specifies deducing oxidising reactivity from E° and hydride thermal stability; this guide focuses on that stated 2027 outcome.
What does the 2027 H2 syllabus require here?
Section 5 requires describing and deducing the relative oxidising reactivity of Group 17 elements using standard electrode potentials.
Continue with connected Chemistry study
- JC1 H2 Chemistry | Hydrogen Halide Thermal Stability Is a Different Group 17 Question
- Hougang H2 Chemistry | Group 17 Oxidising Power and Hydride Trends
- Hougang Chemistry Tuition and Tutor Library
Parents can bring the student’s 2027 H2 syllabus, a standard-potential data table and one incorrect displacement prediction when discussing tuition. Verify teaching groups, venue, class availability and fees directly. This article does not guarantee outcomes or instruct handling hazardous halogens.
Official sources, limitations and safety
- SEAB 2027 H2 Chemistry 9476, Section 5 Periodic Table
- SEAB Chemistry Data Booklet, standard redox potentials
- SEAB 2027 A-Level syllabus directory
- Immutable eduKateSG Clementi teaching reference
- Hougang Chemistry Tuition and Tutor Library
The 2027 H2 9476 syllabus specifically requires deducing Group 17 oxidising strength from E° values. The +1.36, +1.07 and +0.54 V reduction potentials are from the SEAB data booklet. The voltages imply standard-state thermodynamic directions under the listed species conditions, not measured rates or unconditional outcomes at every concentration. The topic is handled as paper redox reasoning, without chlorine, bromine or iodine preparation procedures.