From direct transfer to electrode processes
When a zinc rod is dipped directly into copper sulphate solution, a redox reaction proceeds at once: zinc is oxidised to zinc ions while copper ions are reduced to metallic copper, the electrons passing straight from zinc to the copper ion at the metal surface. NCERT notes that heat is evolved in this direct transfer. The reaction is genuinely a redox process, but the energy is released as heat and is lost for any electrical purpose.
The whole of §7.4 turns on a single modification: carry out the same reaction but force the electron transfer to occur indirectly. To do this, the zinc metal is separated from the copper sulphate solution. Copper sulphate is taken in one beaker with a copper strip, and zinc sulphate in a second beaker with a zinc strip. At the interface of each metal and its salt solution, both the oxidised and the reduced forms of the same species coexist — precisely the species of a reduction or an oxidation half-reaction.
The overall Daniell-cell reaction split into its two halves:
Oxidation (at the zinc electrode): $\ce{Zn(s) -> Zn^2+(aq) + 2e^-}$
Reduction (at the copper electrode): $\ce{Cu^2+(aq) + 2e^- -> Cu(s)}$
Overall: $\ce{Zn(s) + Cu^2+(aq) -> Zn^2+(aq) + Cu(s)}$
Redox couples and the Daniell cell
NCERT defines a redox couple as the oxidised and reduced forms of a substance taking part together in an oxidation or reduction half-reaction. The couple is written with the oxidised form before the reduced form, separated by a slash or vertical line that represents the interface, for example $\ce{Zn^2+}/\ce{Zn}$ and $\ce{Cu^2+}/\ce{Cu}$. Each electrode in the cell is one redox couple.
The two beakers are placed side by side and joined by a salt bridge — a U-tube of potassium chloride or ammonium nitrate solution, solidified with agar-agar into a jelly. The bridge gives electrical contact between the two solutions without letting them mix. The zinc and copper rods are wired together externally through an ammeter and a switch. This complete set-up is the Daniell cell.
With the switch off, nothing happens. The moment the switch closes, NCERT records two observations: the electron transfer no longer goes directly from $\ce{Zn}$ to $\ce{Cu^2+}$ but travels through the external metallic wire, and electricity flows between the two solutions by migration of ions through the salt bridge. The salt bridge thus completes the inner circuit and keeps each half-cell electrically neutral as the reaction proceeds.
Electrode potential and the standard hydrogen electrode
Current flows only because a potential difference exists between the copper and zinc rods, here called electrodes. The potential associated with each electrode is its electrode potential. Physically, it measures the tendency of a metal to lose electrons and pass into solution as ions, or to gain electrons from its ions — the competition between dissolution and deposition at the metal/solution interface.
To compare electrodes fairly, conditions are fixed. When every species in the electrode reaction is at unit concentration (and any gas is confined to one atmosphere), with the reaction at 298 K, the potential is the standard electrode potential, written $E^\circ$. A single electrode's absolute potential cannot be measured; only the difference between two electrodes can. NCERT therefore fixes a reference: by convention the standard electrode potential of the hydrogen electrode is taken as exactly 0.00 V.
"More positive E° = stronger oxidising agent" — and the SHE zero is a convention
Two facts the examiners exploit every year. First, the entries of Table 7.1 are reduction half-reactions, so a more positive $E^\circ$ means the oxidised form is more readily reduced and is therefore a stronger oxidising agent (and the lower the E°, the stronger the reducing agent). Do not invert this. Second, the standard hydrogen electrode is assigned $E^\circ = 0.00\ \text{V}$ purely as a chosen reference, not because hydrogen has "no" potential — every other E° is reported relative to it.
More positive $E^\circ$ → stronger oxidant; more negative $E^\circ$ → stronger reductant; $E^\circ_{\ce{H+}/\ce{H2}} = 0.00\ \text{V}$ by definition.
NCERT also fixes the sign meaning of E° relative to this reference. A negative $E^\circ$ means the redox couple is a stronger reducing agent than the $\ce{H+}/\ce{H2}$ couple; a positive $E^\circ$ means the couple is a weaker reducing agent than $\ce{H+}/\ce{H2}$. The further a value sits from zero, the more lopsided that tendency.
The standard electrode potential series
Tabulating these reference values for many electrodes gives the standard electrode potential series (also met as the electrochemical or activity series). NCERT Table 7.1 lists selected reduction half-reactions at 298 K with their $E^\circ$ values; ions are aqueous, $\ce{H2O}$ is liquid, and gases and solids are marked. Reading down the column, oxidising strength falls and reducing strength rises.
| Reduction half-reaction (oxidised form + ne⁻ → reduced form) | E° / V |
|---|---|
| $\ce{F2(g) + 2e^- -> 2F^-}$ | +2.87 |
| $\ce{Co^3+ + e^- -> Co^2+}$ | +1.81 |
| $\ce{H2O2 + 2H+ + 2e^- -> 2H2O}$ | +1.78 |
| $\ce{MnO4^- + 8H+ + 5e^- -> Mn^2+ + 4H2O}$ | +1.51 |
| $\ce{Au^3+ + 3e^- -> Au(s)}$ | +1.40 |
| $\ce{Cl2(g) + 2e^- -> 2Cl^-}$ | +1.36 |
| $\ce{Cr2O7^2- + 14H+ + 6e^- -> 2Cr^3+ + 7H2O}$ | +1.33 |
| $\ce{O2(g) + 4H+ + 4e^- -> 2H2O}$ | +1.23 |
| $\ce{Br2 + 2e^- -> 2Br^-}$ | +1.09 |
| $\ce{Ag+ + e^- -> Ag(s)}$ | +0.80 |
| $\ce{Fe^3+ + e^- -> Fe^2+}$ | +0.77 |
| $\ce{I2(s) + 2e^- -> 2I^-}$ | +0.54 |
| $\ce{Cu^2+ + 2e^- -> Cu(s)}$ | +0.34 |
| $\ce{2H+ + 2e^- -> H2(g)}$ | 0.00 (reference) |
| $\ce{Pb^2+ + 2e^- -> Pb(s)}$ | −0.13 |
| $\ce{Fe^2+ + 2e^- -> Fe(s)}$ | −0.44 |
| $\ce{Zn^2+ + 2e^- -> Zn(s)}$ | −0.76 |
| $\ce{Al^3+ + 3e^- -> Al(s)}$ | −1.66 |
| $\ce{Mg^2+ + 2e^- -> Mg(s)}$ | −2.36 |
| $\ce{Na+ + e^- -> Na(s)}$ | −2.71 |
| $\ce{K+ + e^- -> K(s)}$ | −2.93 |
| $\ce{Li+ + e^- -> Li(s)}$ | −3.05 |
The two extremes anchor the whole series. $\ce{F2}$ at $+2.87\ \text{V}$ is the strongest oxidising agent in the table; $\ce{Li}$ at $-3.05\ \text{V}$ is the strongest reducing agent. The hydrogen electrode at $0.00\ \text{V}$ sits in the middle as the agreed origin.
Reading oxidising and reducing strength from E°
Because every entry is a reduction, a higher $E^\circ$ signals that the oxidised form pulls electrons more strongly — it is more easily reduced and therefore a more powerful oxidising agent. Fluorine, sitting at the top, is the strongest oxidant; its position is consistent with the chapter's earlier statement that the oxidising power of halogens decreases down the group ($\ce{F2 > Cl2 > Br2 > I2}$), mirrored by the falling E° values $+2.87, +1.36, +1.09, +0.54\ \text{V}$.
The reduced forms tell the opposite story. The lower the $E^\circ$, the more readily the reduced form gives up electrons, so the stronger a reducing agent it is. Lithium, sodium and potassium metals — deep in negative territory — are vigorous reductants, while the noble metals near the top ($\ce{Ag}$, $\ce{Au}$) barely reduce anything. A useful ordering question follows directly: to rank oxidising strength, rank the E° values; the most positive wins.
Electrode potential is only the formal scoring layer over the electron-transfer view of oxidation and reduction. Revisit Redox as Electron Transfer to see why "loss of electrons = oxidation" underlies every E° in Table 7.1.
Predicting spontaneity of redox reactions
The practical payoff of the series is prediction. Any candidate redox reaction is split into a reduction half (cathode) and an oxidation half (anode), and the standard cell potential is computed from the reduction potentials:
$$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$$
If $E^\circ_{\text{cell}} > 0$ the reaction is spontaneous as written; if $E^\circ_{\text{cell}} < 0$ the reverse direction is favoured. Equivalently, the couple with the higher reduction potential takes the electrons (gets reduced) and oxidises the couple with the lower one. NCERT Exercise 7.26 asks exactly this — whether $\ce{Fe^3+}$ will oxidise $\ce{I^-}$, whether $\ce{Ag+}$ will oxidise $\ce{Cu}$, and so on.
Will $\ce{Fe^3+}$ oxidise $\ce{I^-}$ to $\ce{I2}$ under standard conditions? Given $E^\circ_{\ce{Fe^3+}/\ce{Fe^2+}} = +0.77\ \text{V}$ and $E^\circ_{\ce{I2}/\ce{I^-}} = +0.54\ \text{V}$.
Cathode (reduction): $\ce{Fe^3+ + e^- -> Fe^2+}$, $E^\circ = +0.77\ \text{V}$.
Anode (oxidation): $\ce{2I^- -> I2 + 2e^-}$ (the $\ce{I2}/\ce{I^-}$ couple, $E^\circ = +0.54\ \text{V}$).
$E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}} = 0.77 - 0.54 = +0.23\ \text{V}$.
Since $E^\circ_{\text{cell}} > 0$, the reaction $\ce{2Fe^3+ + 2I^- -> 2Fe^2+ + I2}$ is spontaneous: $\ce{Fe^3+}$ does oxidise iodide. Its higher reduction potential confirms $\ce{Fe^3+}$ is the stronger oxidant of the pair.
Redox reactions as the basis for titrations
The same electron-transfer chemistry underlies redox titrations (§7.3.3), where the strength of a reductant or oxidant is found against a standard solution using a redox-sensitive indicator. NCERT illustrates three indicator strategies, summarised below.
| Method | How the end-point is signalled |
|---|---|
| Self-indicator — $\ce{MnO4^-}$ | Intensely coloured permanganate acts as its own indicator. The first lasting tinge of pink (at $\ce{MnO4^-}$ as low as $10^{-6}\ \text{mol L}^{-1}$) appears once all reductant ($\ce{Fe^2+}$ or $\ce{C2O4^2-}$) is oxidised, minimising overshoot beyond the equivalence point. |
| External redox indicator — $\ce{Cr2O7^2-}$ | Dichromate is not a self-indicator. It oxidises the indicator diphenylamine just after the equivalence point, producing an intense blue colour that marks the end-point. |
| Iodometric (starch) | Reagents that oxidise $\ce{I^-}$, e.g. $\ce{2Cu^2+ + 4I^- -> Cu2I2 + I2}$, liberate iodine, which gives an intense blue with starch. The colour vanishes as thiosulphate consumes the iodine: $\ce{I2 + 2S2O3^2- -> 2I^- + S4O6^2-}$. |
Limitation of the oxidation-number concept
NCERT §7.3.4 closes the conceptual arc by noting that the idea of redox keeps evolving. The oxidation-number bookkeeping is powerful but formal — it assigns whole-number charges that atoms do not literally carry. In the most recent view, oxidation is better described as a decrease in electron density around the atom and reduction as an increase in electron density. This electron-density picture connects cleanly to electrode processes, where the genuine driving force is the redistribution of electrons between coupled redox couples rather than any tally of formal oxidation states.
Redox & electrode processes in one glance
- Separating the oxidation and reduction halves of a redox reaction (e.g. $\ce{Zn + Cu^2+}$) routes electrons through a wire — the Daniell cell.
- A redox couple = oxidised form / reduced form ($\ce{Zn^2+}/\ce{Zn}$); each electrode is one couple.
- Standard electrode potential $E^\circ$ is measured at unit concentration, 1 atm gas, 298 K, against the SHE fixed at $0.00\ \text{V}$ by convention.
- Table 7.1 lists reduction half-reactions: more positive $E^\circ$ → stronger oxidant ($\ce{F2}$, $+2.87\ \text{V}$); more negative → stronger reductant ($\ce{Li}$, $-3.05\ \text{V}$).
- $E^\circ_{\text{cell}} = E^\circ_{\text{cathode}} - E^\circ_{\text{anode}}$; positive value → spontaneous reaction.
- Redox titrations use self-indicators ($\ce{MnO4^-}$), diphenylamine with $\ce{Cr2O7^2-}$, or the starch–iodine–thiosulphate system.