Q.Value of standard electrode potential for the oxidation of $Cl^-$ ions is more positive than that of water, even then in the electrolysis of aqueous sodium chloride, why is $Cl^-$ oxidised at anode instead of water?

The key idea is overpotential: the actual potential required to oxidise water is much higher than its thermodynamic value due to a kinetic barrier at the electrode surface.

Reasoning:

1. The standard electrode potential for $Cl^-$ oxidation is $E^\circ = +1.36\ \text{V}$ (for $2Cl^- \rightarrow Cl_2 + 2e^-$), while for water oxidation it is $E^\circ = +1.23\ \text{V}$ (for $2H_2O \rightarrow O_2 + 4H^+ + 4e^-$). Thermodynamically, water should oxidise first.

2. However, the oxidation of water at an inert electrode (like platinum or graphite) requires a large overpotential — often $0.4$–$0.6\ \text{V}$ extra — because the reaction involves multiple steps and slow electron transfer. …

The key here is that standard electrode potentials are thermodynamic values — they tell you which reaction is spontaneous under standard conditions, but they don't account for the kinetic barrier (overpotential) that many electrode reactions face.

Let’s look at the two possible oxidation reactions at the anode:

1. Oxidation of Cl⁻ (to chlorine gas):

2. Oxidation of water (to oxygen gas):

Thermodynamically, water oxidation has a less positive (more favourable) potential — it should happen first. But in practice, the opposite occurs. Here’s why.

Step-by-step reasoning

1. Standard potentials are equilibrium values

The $E^\circ$ values above are measured under ideal conditions (1 M concentration, 1 atm pressure, smooth platinum electrode). They tell you the voltage at which the reaction just begins to occur. But real electrolysis happens at higher voltages because of overpotential — the extra voltage needed to overcome the activation energy barrier.

2. Oxygen evolution has a large overpotential on common anodes

On inert electrodes like platinum or graphite, the reaction:

requires a significant overpotential — often around 0.4–0.6 V. This means the actual potential needed to evolve oxygen is:

3. Chlorine evolution has a much smaller overpotential

The reaction:

has a very low overpotential on platinum or graphite — often less than 0.1 V. So its actual potential is close to the thermodynamic value:

4. Compare the actual potentials

Now the effective potentials are:

• Cl⁻ oxidation: ≈ –1.36 V
• Water oxidation: ≈ –1.6 to –1.8 V

The reaction requiring a less negative (i.e., less energy-demanding) potential will occur first. Cl⁻ oxidation now has the advantage. …

Concept: Electrode Kinetics vs Thermodynamics (Overpotential)

The key idea is that thermodynamic favourability (standard electrode potential) does not always determine which reaction occurs — kinetic factors like overpotential can reverse the order.

Method: Overpotential Analysis

Step 1: Write the two possible oxidation reactions at the anode

• Oxidation of chloride ions:

• Oxidation of water:

Thermodynamically, water oxidation has a less positive $E^\circ$, so it should occur first. But in practice, chlorine is produced.

Step 2: Introduce the concept of overpotential

• Overpotential ($\eta$) is the extra voltage required beyond the thermodynamic potential to drive a reaction at a noticeable rate.
• For oxygen evolution (water oxidation), the overpotential on common anode materials (like platinum or graphite) is very high — often $>0.6 \, \text{V}$.
• For chlorine evolution, the overpotential is much smaller (typically $<0.1 \, \text{V}$).

Step 3: Compare the actual required potentials

• Actual potential needed for $Cl^-$ oxidation:

• Actual potential needed for water oxidation: …

🧠 The Core Concept

In the electrolysis of aqueous NaCl, two species can be oxidised at the anode:

• Cl⁻ ions (from NaCl)
• H₂O molecules (from the solvent)

The standard electrode potentials (at 25°C, 1 M, 1 atm) are:

• For oxidation of Cl⁻:

• For oxidation of water:

Key point: A more positive (or less negative) $E^\circ$ means the reduction is easier — but for oxidation, we flip the sign. So:

• Oxidation of water: $E^\circ_{\text{ox}} = -1.23 \, \text{V}$
• Oxidation of Cl⁻: $E^\circ_{\text{ox}} = -1.36 \, \text{V}$

Since $-1.23 > -1.36$, water should be oxidised more easily — yet in practice, Cl⁻ is oxidised at the anode. Why?

✗ Common Mistake #1: Ignoring Overpotential

The error: Students compare only standard potentials and conclude water must oxidise.

Why it's wrong: Standard potentials assume reversible conditions. In real electrolysis, overpotential (extra voltage needed to overcome kinetic barriers) matters — especially for gas evolution.

• O₂ evolution from water has a high overpotential on common anode materials (like graphite or platinum).
• Cl₂ evolution has a much lower overpotential.

So the actual voltage required for water oxidation becomes more negative than for Cl⁻ oxidation under real conditions.

How to avoid: Always remember — in electrolysis, the actual ease of oxidation depends on:

The species with the less negative (or more positive) actual potential gets oxidised.

✗ Common Mistake #2: Confusing Reduction vs Oxidation Potentials

The error: Students see $E^\circ(\text{Cl}_2/\text{Cl}^-) = +1.36 \, \text{V}$ and $E^\circ(\text{O}_2/\text{H}_2\text{O}) = +1.23 \, \text{V}$ and think Cl⁻ is harder to oxidise because its reduction potential is higher.

Why it's wrong: These are reduction potentials. For oxidation, you must reverse the sign:

• Oxidation of Cl⁻: $E^\circ_{\text{ox}} = -1.36 \, \text{V}$
• Oxidation of water: $E^\circ_{\text{ox}} = -1.23 \, \text{V}$

Now compare: $-1.23 > -1.36$, so water is thermodynamically easier to oxidise.

How to avoid: Always write the half-reaction in the direction it occurs (oxidation at anode) and use the corresponding potential. If given reduction potentials, flip the sign for oxidation.

✗ Common Mistake #3: Forgetting Concentration Effects

The error: Students assume standard conditions (1 M Cl⁻) — but in brine electrolysis, concentrated NaCl is used.

Why it matters: The Nernst equation shows that higher [Cl⁻] makes its oxidation easier (less negative potential): …