Redox reactions and electrode processes

What are electrode processes?

In plain words: drop zinc into copper sulphate solution and the electrons jump straight from zinc atoms to copper ions, so no useful external electrical work is collected; the released energy largely heats the surroundings. Put the zinc and the copper ions in separate beakers, join them with a wire and an ionic connection, and the electrons flow through the wire to get there: the reaction now drives an electric current. How strongly each couple pushes or pulls electrons is measured by its electrode potential.

From direct to indirect electron transfer

When a zinc rod is dipped in copper sulphate solution, zinc is oxidised to zinc ions and copper ions are reduced to copper by direct transfer of electrons, and heat is evolved (NCERT pages 249–250). To make the transfer indirect, the zinc must be kept away from the copper sulphate solution:

• one beaker holds copper sulphate solution with a copper strip;
• the other holds zinc sulphate solution with a zinc strip.

At the surface of each metal, both the oxidised and the reduced forms of the same species are present, Cu²⁺ and Cu in one beaker, Zn²⁺ and Zn in the other. These are the two redox couples. The slash here names a redox couple; in full cell notation a vertical line marks a phase boundary, and the sequence follows the cell layout. Couples are conventionally named with the oxidised form first: Zn²⁺/Zn and Cu²⁺/Cu (page 250).

The Daniell cell

The two beakers stand side by side, and their solutions are joined by a salt bridge. This is a U-tube holding a solution of potassium chloride or ammonium nitrate, usually set to a jelly by boiling with agar-agar and cooling. It connects the solutions electrically while minimising bulk mixing. The zinc and copper rods are joined by a metal wire through an ammeter and, in NCERT’s set-up, a switch (NCERT Fig. 7.3, page 250). This set-up is the Daniell cell.

• With the switch off, no sustained external current flows; interfacial charge equilibration and possible side reactions need not stop.
• With the switch on, electrons pass from Zn to Cu²⁺ not directly but through the wire, and the ammeter shows a current.
• Inside the cell, the circuit is completed by ions moving through the salt bridge: anions move into the zinc beaker, where Zn²⁺ ions are forming, and cations into the copper beaker, where Cu²⁺ ions are being removed, so each solution stays electrically neutral. Potassium chloride and ammonium nitrate are strong electrolytes, so the bridge is full of mobile ions.

Zinc is oxidised at the zinc rod, the anode, Zn → Zn²⁺ + 2e⁻, which therefore becomes the negative electrode. Copper ions take up the electrons at the copper rod, the cathode, Cu²⁺ + 2e⁻ → Cu, the positive electrode. The electrons flow through the wire from zinc to copper; the conventional current flows the opposite way, from copper to zinc (page 250).

Electrode potentials

A current can flow only if there is a potential difference between the two rods, now called electrodes. The potential of each electrode is its electrode potential (page 250). The standard electrode potential, E^⊖, refers to unit activities and a specified temperature, commonly 298 K in tables. For an ideal dilute solute, activity is approximated by concentration divided by its standard concentration; a gas uses its 1 bar standard state, and a pure solid or liquid has unit activity. NCERT uses the older 1 atm reference; the pressure-reference correction is small for the examples here but depends on the half reaction.

By convention, the standard hydrogen electrode, 2H⁺ + 2e⁻ → H₂(g), has E^⊖ = 0.00 V. Each E^⊖ measures the tendency of the couple to stay in its oxidised or its reduced form (pages 250–251):

• a negative E^⊖ means its reduced form is a stronger reducing agent than H₂ under standard conditions;
• a positive E^⊖ means its oxidised form is a stronger oxidising agent than H⁺ under standard conditions; correspondingly its reduced form is a weaker reducing agent than H₂.

The table of standard electrode potentials

NCERT Table 7.1 (page 251) lists reductions, oxidised form + ne⁻ → reduced form, at 298 K. Some of its values:

Going up the table, the oxidised forms on the left become stronger oxidising agents, with fluorine the strongest among these listed couples. Going down, the reduced forms become stronger reducing agents, with lithium the strongest among those listed under the specified aqueous standard conditions (page 251).

Using the table. A reaction between the oxidised form of one couple and the reduced form of another has a favourable standard-state Gibbs change if the first couple has the higher E^⊖. Actual reaction direction follows the actual cell potential, which depends on activities; surface films, kinetics and competing reactions also matter. Copper (Cu²⁺/Cu, +0.34 V) is oxidised by Ag⁺ (+0.80 V) but not by H⁺ (0.00 V), which is why copper does not give hydrogen with dilute hydrochloric acid while zinc (−0.76 V) does. This is the same order of electron-releasing tendency, Zn > Cu > Ag, that NCERT Section 7.2.1 found from displacement reactions. Note that in this table the best reducing agents are at the bottom.

In Class 12 you will find the voltage of a cell from the two potentials, E^⊖_cell = E^⊖_cathode − E^⊖_anode. For the Daniell cell this is 0.34 − (−0.76) = 1.10 V. Both values are reduction potentials read straight from the table; do not change the sign of the anode’s value before subtracting. A positive E^⊖_cell means ΔG° is negative. In Example2(d), 0.77 − 0.80 = −0.03 V makes ΔG° positive, but the reaction can still proceed from a mixture with sufficiently low product activities.

Solved examples

Common mistakes to avoid

Quick revision summary

• Redox couple: oxidised and reduced forms together, written oxidised/reduced (Zn²⁺/Zn).
• Daniell cell: a zinc rod in ZnSO₄ and a copper rod in CuSO₄ joined by a wire and a salt bridge; anode Zn (−), cathode Cu (+); electrons from Zn to Cu through the wire.
• Standard electrode potential: unit activities, gas standard pressure 1 bar (NCERT: 1 atm), specified temperature (table: 298 K); H⁺/H₂ = 0.00 V by convention.
• More positive E^⊖: stronger oxidising agent (F₂); more negative: stronger reducing agent (Li).
• The oxidant’s couple having higher E^⊖ gives favourable standard-state driving force; actual direction depends on activities.

Questions students ask

Is a Daniell cell the same as a galvanic cell?

A Daniell cell is one particular galvanic cell, the zinc–copper cell Zn | Zn²⁺(aq) || Cu²⁺(aq) | Cu, while galvanic cell (also called voltaic cell) is the general name for any cell in which a spontaneous redox reaction produces electrical energy. So every Daniell cell is a galvanic cell, but not every galvanic cell is a Daniell cell: compatible half-cells with different actual electrode potentials, such as Pb²⁺/Pb and Ag⁺/Ag, can make one when a wire and suitable ionic connection complete the circuit. The opposite device is an electrolytic cell, in which an external supply drives a non-spontaneous redox reaction, as in the electrolysis of molten sodium chloride.

Why is the standard electrode potential of the hydrogen electrode zero?

The standard hydrogen electrode has E^⊖ = 0.00 V by convention, not by measurement, because the potential of a single electrode cannot be measured on its own: a voltmeter can only measure the difference between two electrodes. Chemists therefore chose one reference, platinum in contact with H₂ gas at 1 bar and H⁺ ions at unit activity, assigned it zero at every temperature, and quote every other standard electrode potential as the voltage of a cell made with it. So Al³⁺/Al at −1.66 V means that a standard aluminium electrode is 1.66 V negative with respect to the hydrogen electrode.

Why is KCl or KNO₃ used in a salt bridge?

A salt bridge is filled with KCl, KNO₃ or NH₄NO₃ because the cation and the anion of each salt move at almost the same speed, so they carry nearly equal shares of the current and the extra potential that forms where two solutions meet, the liquid junction potential, stays very small. The salt must also not react with either solution. KCl is therefore avoided when a beaker contains Ag⁺, Pb²⁺ or Hg₂²⁺ ions, which would precipitate as chlorides, and KNO₃ or NH₄NO₃ is used instead.

What is the difference between electrode potential and cell potential?

Electrode potential belongs to a single electrode, one couple such as Cu²⁺/Cu, and is measured against the standard hydrogen electrode, while cell potential, or emf, is the difference between the potentials of the two electrodes of a cell when no current is drawn: E^⊖_cell = E^⊖_cathode − E^⊖_anode. For a cell made from Mg²⁺/Mg (−2.36 V) and Cu²⁺/Cu (+0.34 V), magnesium is the anode and E^⊖_cell = 0.34 − (−2.36) = 2.70 V. How the emf changes with concentration, through the Nernst equation, is Class 12 Electrochemistry.

What is the difference between the activity series and the electrochemical series?

In NCERT Class 11 the activity series and the electrochemical series are one list under two names, but many books separate them: the activity series ranks metals by how vigorously they react, and the electrochemical series ranks couples of metals and non-metals by standard electrode potential, from F₂/F⁻ to Li⁺/Li. The orders mostly agree, as a lower E^⊖ means a stronger reducing agent. They differ where speed matters: lithium has the most negative E^⊖ yet reacts with water more gently than sodium, as E^⊖ shows how far a reaction goes, not how fast.

Is electrochemistry in Class 11 or Class 12?

Electrochemistry is mainly a Class 12 chapter: NCERT Class 12 Chemistry Part I, Unit 2, covers galvanic and electrolytic cells, the Nernst equation, conductance, batteries and fuel cells. Class 11 introduces it at the end of Redox Reactions, in Section 7.4, with redox couples, the Daniell cell and standard electrode potentials. The JEE Main 2027 and NEET (UG) 2026 syllabuses join both parts in one unit, Unit 7: Redox Reactions and Electrochemistry, so Class 11 redox and Class 12 electrochemistry are best revised together.

Practice questions

Try each question first, then open the answer to check your working.

1. Q1. Which is the stronger reducing agent, zinc or iron? Use NCERT Table 7.1.

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   Zinc: Zn²⁺/Zn is −0.76 V, more negative than Fe²⁺/Fe, −0.44 V, so zinc gives up electrons more readily.

2. Q2. Why does zinc give hydrogen with dilute hydrochloric acid, but copper does not?

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   H⁺/H₂ is 0.00 V. Zn²⁺/Zn (−0.76 V) is below it, so zinc can reduce H⁺ to H₂. Cu²⁺/Cu (+0.34 V) is above it, so copper cannot. Copper does dissolve in nitric acid, but there the oxidant is nitrate, NO₃⁻/NO (+0.97 V in NCERT Table 7.1), not H⁺, and no hydrogen is given off.

3. Q3. In the Daniell cell, what is the job of the salt bridge, and what happens if it is removed?

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   It completes the circuit inside the cell by the movement of ions, without letting the solutions mix, and keeps each solution electrically neutral. Without it, charge builds up in the beakers, sustained current rapidly ceases and the net cell reaction is arrested; a brief charging transient can occur.