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Will an iron nail turn copper sulphate solution colourless? Does copper dissolve in hydrochloric acid? Can iodine free bromine from bromide? Each looks like a fact to memorise, but one table and one subtraction answer all of them. This post shows the rule, works seven predictions, and explains where the method stops working.
What is the electrochemical series?
The NCERT Class 11 unit on redox reactions introduces it as an "activity series or electrochemical series": zinc releases electrons to copper, copper to silver, so Zn > Cu > Ag. The NCERT Class 12 unit on electrochemistry puts numbers on it in Table 2.1, standard electrode potentials at 298 K, written as reductions.
NCERT's reading of the table: fluorine, at +2.87 V, is the strongest oxidising agent; lithium, at −3.05 V, is the strongest reducing agent in aqueous solution. Going down the table, the oxidising power of the species on the left falls and the reducing power of the species on the right rises.
The JEE Advanced 2026 syllabus names "Electrochemical series, emf of galvanic cells", and the JEE (Main) 2026 syllabus covers standard electrode potentials and cell emf. Check the current year's syllabus. For cell notation and the Nernst equation, see our post on cell EMF and the Nernst equation; this one is about prediction.
The rule
For a proposed reaction, decide which species is reduced and which is oxidised, then:
E°cell = E°(couple being reduced) − E°(couple being oxidised)
NCERT gives ΔG = −nFE(cell). A positive E°cell makes ΔG° negative, so the reaction is feasible under standard conditions. A negative E°cell means the reverse reaction is the feasible one.
Seven worked predictions
Values are from NCERT Table 2.1.
Iron nail in copper sulphate. Cu²⁺ is reduced (+0.34 V), Fe is oxidised to Fe²⁺ (−0.44 V). E° = 0.34 − (−0.44) = +0.78 V. Yes: copper deposits on the nail.
Storing copper sulphate in a zinc pot (an NCERT in-text question). Cu²⁺ reduced, Zn oxidised (−0.76 V). E° = 0.34 − (−0.76) = +1.10 V. The zinc would react, so no.
Copper in hydrochloric acid. H⁺ reduced (0.00 V), Cu oxidised (+0.34 V). E° = 0.00 − 0.34 = −0.34 V. No reaction.
Copper in nitric acid. Nitrate is reduced to NO (+0.97 V). E° = 0.97 − 0.34 = +0.63 V. Yes. Copper resists HCl but dissolves in nitric acid because nitrate, not H⁺, is the oxidising agent.
Chlorine with bromide. Cl₂ reduced (+1.36 V), Br⁻ oxidised (+1.09 V). E° = 1.36 − 1.09 = +0.27 V. Yes, which matches NCERT's Class 11 equation Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂.
Iodine with bromide. I₂ reduced (+0.54 V), Br⁻ oxidised (+1.09 V). E° = 0.54 − 1.09 = −0.55 V. No.
Iron(III) with iodide and with bromide. Fe³⁺/Fe²⁺ is +0.77 V. With I⁻: 0.77 − 0.54 = +0.23 V, yes. With Br⁻: 0.77 − 1.09 = −0.32 V, no.
A shortcut follows: a species on the left of the table can oxidise any species on the right that sits below it. So any couple above +0.77 V, such as Br₂, Cr₂O₇²⁻, Cl₂ or MnO₄⁻, can oxidise Fe²⁺ to Fe³⁺, which answers another NCERT in-text question.
Where does the method stop working?
It predicts feasibility, not speed. A positive E° says the reaction can go, not that it goes quickly.
It assumes standard conditions. E° values are for 1 M solutions at 298 K. When E° is small, concentration can tip the balance. Ag⁺ with Fe²⁺ gives 0.80 − 0.77 = +0.03 V, a margin the Nernst equation can wipe out.
Water can react first. Sodium sits far below copper, so you might predict that sodium displaces copper from copper sulphate solution. But NCERT's Class 11 unit shows that alkali metals displace hydrogen from cold water, 2Na + 2H₂O → 2NaOH + H₂, and that is what happens.
Research shows why these limits matter. In a study of college students, Sanger and Greenbowe (1997) found misconceptions that half-cell potentials are absolute and can predict the spontaneity of a single half-cell, and that cell potentials do not depend on concentration. Most of those students could still calculate cell potentials correctly. Getting the number right does not mean the idea is clear.
Which mistakes cost marks?
Reversing the subtraction. Write which species is reduced before you subtract.
Using a half-cell on its own. A reaction needs one reduction and one oxidation; a single E° predicts nothing.
Multiplying E° by coefficients.
Ignoring water when a very reactive metal is in solution.
Calling a reaction fast because E° is large.
A practice routine
From Table 2.1, pick ten pairs and predict each reaction with the subtraction rule.
For three of them, write the balanced equation; our post on balancing redox reactions has the half-reaction method.
Explain one case where E° misleads, using concentration or water.
Apply the table to rusting: our post on corrosion as electrochemistry shows how.
If the signs keep tripping you up, the Fermi Problem Library has worked problems with a step-by-step solver that explains the reasoning rather than printing an answer.
Tonight, predict whether zinc dissolves in dilute hydrochloric acid, write the subtraction, and check your answer against NCERT's Class 11 examples.
