Electrochemistry 05 Electrolysis CUET - CUET Accountancy

Electrolysis and Faraday’s Laws

Electrolytic Cell: Structure and Function

An electrolytic cell is a device in which non-spontaneous chemical reactions are driven by the application of external electrical energy. In these cells, electrical energy is converted into chemical energy, and a battery is used to provide the necessary external potential. The cell contains two electrodes (anode and cathode) immersed in an electrolyte solution.

• Non-spontaneous reactions occur; external energy is required.
• Electrical energy is converted to chemical energy.
• Battery provides the external potential to drive the reaction.
• Anode is positively charged; cathode is negatively charged.…

Electrolysis and Faraday’s Laws

Electrolytic Cell: Structure and Function

An electrolytic cell is a device in which non-spontaneous chemical reactions are driven by the application of external electrical energy. In these cells, electrical energy is converted into chemical energy, and a battery is used to provide the necessary external potential. The cell contains two electrodes (anode and cathode) immersed in an electrolyte solution.

• Non-spontaneous reactions occur; external energy is required.
• Electrical energy is converted to chemical energy.
• Battery provides the external potential to drive the reaction.
• Anode is positively charged; cathode is negatively charged.
• Example: In molten NaCl, $Na^+$ migrates to the cathode (reduction), $Cl^-$ migrates to the anode (oxidation).

Electrode Reactions and Identification

In an electrolytic cell, oxidation and reduction occur at the electrodes. The identification of anode and cathode is based on the type of reaction occurring at each electrode.

• Anode: Site of oxidation (loss of electrons), positively charged.
• Cathode: Site of reduction (gain of electrons), negatively charged.
• At the cathode, cations are reduced; at the anode, anions are oxidized.
• Example reactions:
  • At cathode: $Na^+ + e^- \rightarrow Na(s)$
  • At anode: $2Cl^- \rightarrow Cl_2(g) + 2e^-$

Electrolysis: Definition and Key Features

Electrolysis is the process in which electrical energy is used to drive non-spontaneous chemical reactions. It always takes place in an electrolytic cell and involves the movement of ions towards oppositely charged electrodes.

• Electrolysis: Decomposition of a substance by passing electric current.
• Only non-spontaneous reactions are made to occur by applying an external battery.
• At the cathode, cations are reduced; at the anode, anions are oxidized.
• The negative electrode is always the cathode; the positive electrode is always the anode.

Calculation of Amounts Deposited or Liberated

The amount of substance deposited or liberated at an electrode during electrolysis is directly related to the number of moles of electrons (or Faradays) passed.

• 1 mole of electrons carries a charge of $96500$ coulombs ($1$ Faraday).
• The number of moles of electrons used equals the number of Faradays passed.
• For a reaction: $Na^+ + e^- \rightarrow Na(s)$, $1$ mole of electrons deposits $1$ mole of sodium.
• For $Cu^{2+} + 2e^- \rightarrow Cu(s)$, $2$ moles of electrons deposit $1$ mole of copper.

Faraday’s Laws of Electrolysis

Faraday’s First Law states that the mass of a substance altered at an electrode is directly proportional to the quantity of electricity passed through the electrolyte.

• First Law: $w \propto Q$ (where $w$ is mass altered, $Q$ is charge passed).
• $w = Z \cdot Q$, where $Z$ is the electrochemical equivalent.
• $Q = I \cdot t$ (current in amperes, time in seconds).
• $Z$ is defined as the mass altered when $1$ ampere current is passed for $1$ second.

Calculation Methods and Key Formulas

To solve electrolysis problems, a systematic approach is used involving calculation of charge, Faradays, moles of electrons, and then relating to the chemical reaction.

1. Calculate charge: $Q = I \cdot t$
2. Find Faradays: Number of Faradays $= \frac{Q}{96500}$
3. Relate Faradays to moles of electrons: Number of Faradays $=$ moles of electrons used
4. Use reaction stoichiometry to find moles/mass/volume of product
5. If current efficiency is given: Effective current $= \text{Efficiency} \times \text{Applied current}$

• For gases at NTP: Volume $= \text{moles} \times 22.4$ L

Preferential Discharge and Electrochemical Series

When more than one cation or anion is present, the ion with higher reduction potential (for cations) or lower oxidation potential (for anions) is preferentially discharged.

• At the cathode, the ion with higher reduction potential is reduced.
• At the anode, the ion with lower oxidation potential is oxidized.
• The electrochemical series is used to compare reduction potentials.
• Example: In aqueous NaCl, $H^+$ is reduced at the cathode (producing $H_2$), $Cl^-$ is oxidized at the anode (producing $Cl_2$).

Special Cases: Nature of Electrodes

The nature of the electrode can affect the products of electrolysis.

• Inert electrodes (Pt, Pd, C): Do not participate in the reaction; only ions are involved.
• Reactive electrodes (e.g., Cu in $CuSO_4$ solution): The electrode itself can undergo oxidation or reduction.
  • If the electrode is the same as the ion in solution, both oxidation and reduction involve the electrode metal.
• Mercury electrode: In presence of $Na^+$, sodium forms an amalgam with mercury at the cathode, regardless of the presence of $H^+$.

Application to Typical Problems

A stepwise approach is used to solve quantitative electrolysis problems:

1. Calculate total charge using current and time.
2. Convert charge to Faradays.
3. Relate Faradays to moles of electrons.
4. Use reaction stoichiometry to determine moles/mass/volume of product.
5. Adjust for current efficiency if specified.

• Example: Passing $96.5$ A for $100$ s gives $9650$ C; number of Faradays $= 0.1$.
• For $Na^+$ reduction: $0.1$ mol electrons deposit $0.1$ mol Na; mass $= 0.1 \times 23 = 2.3$ g.
• For $Cl_2$ evolution: $1$ mol electrons gives $0.5$ mol $Cl_2$; $0.1$ mol electrons gives $0.05$ mol $Cl_2$; volume $= 0.05 \times 22.4 = 1.12$ L.

Effect of Current Efficiency

Current efficiency accounts for the fraction of applied current actually used in the desired electrochemical reaction.

• Effective current $= \text{Applied current} \times \frac{\text{Efficiency}}{100}$
• All subsequent calculations use the effective current.

Exam-Focused Traps and Conceptual Points

• At the cathode, only reduction occurs; at the anode, only oxidation occurs.
• In aqueous solutions, metals above $H^+$ in the electrochemical series cannot be deposited from their aqueous salt solutions (e.g., $Mg^{2+}$).
• The nature of the electrode (inert vs. reactive) can change the products.
• For preferential discharge, always compare reduction potentials for cations and oxidation potentials for anions.
• The medium (acidic, basic, neutral) after electrolysis depends on which ions remain in solution.

Summary Table: Preferential Discharge (as per transcript)

• For cations, the lower the position in the electrochemical series, the more likely it is to be reduced.
• For anions, the transcript specifies $Cl^-$ is discharged over $OH^-$, and $CH_3COO^-$ over $OH^-$.

Key Definitions

• Electrolysis: Decomposition of a compound by passing electric current.

• Electrolytic cell: Device for carrying out electrolysis.

• Faraday (F): $96500$ coulombs, the charge carried by $1$ mole of electrons.

• Electrochemical equivalent (Z): Mass of substance altered by $1$ coulomb of charge.

• Current efficiency: Percentage of applied current actually used in the desired reaction.

• Always follow the stepwise calculation method for quantitative problems.

• Use the electrochemical series for predicting products in competitive discharge scenarios.

• Adjust calculations for current efficiency when specified.