Electrolytic Cells & Electrolysis | Electrochemistry Class 12

Electrolytic Cells & Electrolysis

Non-Spontaneous Redox Reactions | Electrochemistry

1. What is an Electrolytic Cell?

                    Definition: A device that uses electrical energy to drive a non-spontaneous chemical reaction ($\Delta G > 0$).
                

Key Components & Sign Convention

Anode (Positive): Electrode connected to the positive terminal of the battery. Oxidation (Loss of electrons) occurs here.
Cathode (Negative): Electrode connected to the negative terminal of the battery. Reduction (Gain of electrons) occurs here.
Electrolyte: Molten salt or aqueous solution containing ions.

Note: This is opposite to Galvanic cells where Anode is Negative.

2. Preferential Discharge Theory

When multiple ions compete to discharge at an electrode, the ion with the Lower Discharge Potential is discharged first.

A. At Cathode (Reduction)

The cation with the Higher Standard Reduction Potential ($E^\circ_{red}$) is reduced first. (Easier to reduce).

Order of Discharge (Decreasing Ease):
$Au^{3+} > Ag^+ > Cu^{2+} > H^+ > Pb^{2+} > Zn^{2+} > Al^{3+} > Mg^{2+} > Na^+ > K^+$

Rule of Thumb: In aqueous solution, active metals (Group 1, 2, Al) are NOT discharged; Water is reduced to $H_2$ gas instead.

B. At Anode (Oxidation)

The anion with the Lower Reduction Potential (or Higher Oxidation Potential) is oxidized first.

Order of Discharge (Decreasing Ease):
$I^- > Br^- > Cl^- > OH^- > NO_3^- > SO_4^{2-} > F^-$

Rule of Thumb: Stable anions ($SO_4^{2-}, NO_3^-$) are NOT discharged; Water is oxidized to $O_2$ gas instead. However, for halogens ($Cl^-$), see "Overpotential".

3. Important Case Studies

Case 1: Molten NaCl (Platinum Electrodes)

Only ions present: $Na^+, Cl^-$.

Cathode (-): $Na^+ + e^- \rightarrow Na(s)$
Anode (+): $2Cl^- \rightarrow Cl_2(g) + 2e^-$

Case 2: Aqueous NaCl (Concentrated)

Ions present: $Na^+, Cl^-, H^+, OH^-$ (from water).

Cathode (-): Competition between $Na^+$ and $H_2O$. $H_2O$ has higher $E^\circ_{red}$ than $Na^+$.
$2H_2O + 2e^- \rightarrow H_2(g) + 2OH^-$ (Hydrogen gas).
Anode (+): Competition between $Cl^-$ and $H_2O$. Normally $H_2O \to O_2$ is favored, but due to Overpotential of Oxygen, $Cl^-$ oxidizes.
$2Cl^- \rightarrow Cl_2(g) + 2e^-$ (Chlorine gas).

Result: $NaOH$ remains in solution.

Case 3: Aqueous $CuSO_4$ (Copper Electrodes)

Here, the anode is Active (participates in reaction).

Cathode (-): $Cu^{2+} + 2e^- \rightarrow Cu(s)$ (Copper deposits).
Anode (+): Copper anode itself oxidizes because oxidation of Cu ($E^\circ_{ox} = -0.34V$) is easier than oxidation of water ($E^\circ_{ox} = -1.23V$) or sulphate.
$Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-$ (Anode dissolves).

Use: Electro-refining of Copper.

4. Applications

Electroplating: Coating a cheap metal with a noble metal (e.g., Ag on spoon). Object is Cathode, Pure Metal is Anode.
Extraction of Metals: Electrometallurgy for reactive metals like Na, Mg, Al (Hall-Heroult Process).
Purification: Electro-refining (Impure Anode $\to$ Pure Cathode).

Practice Quiz

Test your ability to predict Electrolysis Products.

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