Electrochemistry

Unit 3: Electrochemistry based on the provided image:

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Unit 3: Electrochemistry – Notes

Introduction

• Electrochemistry is the branch of chemistry that deals with:
  • The production of electricity from energy released during spontaneous chemical reactions, and
  • The use of electrical energy to drive non-spontaneous chemical reactions.

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Importance of Electrochemistry

• Plays a crucial role in both theoretical and practical aspects of chemistry.
• Widely used in the production of metals, sodium hydroxide, chlorine, fluorine, and other chemicals through electrochemical methods.

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Applications

• Batteries and fuel cells:
  • Convert chemical energy into electrical energy.
  • Used extensively in instruments and devices.
• Eco-friendly technology:
  • Electrochemical reactions are energy efficient and less polluting.
  • Useful for developing green technologies.
• Biological relevance:
  • Sensory signals from cells to the brain and vice versa, and
  • Cellular communication have electrochemical origins.

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Conclusion

• Electrochemistry is a vast and interdisciplinary field.
• In this unit, only the elementary aspects of electrochemistry will be covered.

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⚡ What is Electrochemistry?

• Electrochemistry is the study of production of electricity from energy released during spontaneous chemical reactions and the use of electrical energy to bring about non-spontaneous chemical transformations.

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🔄 Electrochemical Reactions

• Reactions where redox (oxidation-reduction) processes are involved.

• Example: Zinc metal dipped in copper sulfate solution:

  • Zinc displaces copper.

  • Zn → Zn²⁺ + 2e⁻ (oxidation)

  • Cu²⁺ + 2e⁻ → Cu (reduction)

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🔋 Galvanic Cell (Voltaic Cell)

• A device that converts chemical energy into electrical energy using redox reactions.

• In the Zn-Cu system:

  • Zinc electrode (anode) undergoes oxidation.

  • Copper electrode (cathode) undergoes reduction.

  • Electrons flow from zinc to copper via an external circuit.

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🧪 Electrolyte and Salt Bridge

• The solution containing ions allows charge flow.

• A salt bridge connects the two half-cells:

  • Maintains electrical neutrality.

  • Prevents mixing of solutions.

  • Completes the circuit.

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🔁 Spontaneous vs Non-Spontaneous Reactions

• Spontaneous reaction: Occurs on its own; produces electricity (e.g., galvanic cell).

• Non-spontaneous reaction: Requires external electrical energy to occur (e.g., electrolysis).

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💡 Applications of Electrochemistry

• Used in:

• Electroplating

• Electrorefining

• Electrolytic production of metals (like Al, Na)

• Batteries and fuel cells

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📘 Chapter 3: Electrochemistry

🧾 Page 2 – Summary Notes (NCERT)

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🔬 Electrochemical Cell Setup (Zinc-Copper Example)

• Cell Representation:
  Zn (s) | Zn²⁺ (aq) || Cu²⁺ (aq) | Cu (s)

• Anode (Negative Electrode):

  • Oxidation occurs.

  • Zn (s) → Zn²⁺ (aq) + 2e⁻

• Cathode (Positive Electrode):

  • Reduction occurs.

  • Cu²⁺ (aq) + 2e⁻ → Cu (s)

• Electron Flow:

  • From zinc to copper through the external wire.

• Ion Movement Through Salt Bridge:

  • Maintains electrical neutrality.

  • Cations migrate toward the cathode compartment.

  • Anions migrate toward the anode compartment.

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⚙️ Functions of the Salt Bridge

1. Completes the circuit allowing ion flow.

2. Maintains charge balance by:

   • Allowing migration of ions.

   • Preventing charge accumulation in half-cells.

3. Prevents mixing of different solutions that might otherwise react directly.

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📊 Cell Potential (EMF)

• The electromotive force (emf) is the potential difference between two electrodes in a cell when no current is flowing.

• Denoted as E_cell or simply emf.

• It determines how much electrical work the cell can perform.

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📐 Measurement of Cell Potential

• The cell potential depends on:

  • Nature of electrodes

  • Ion concentrations

  • Temperature

• Measured using a potentiometer or voltmeter when no current flows (open circuit).

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🔁 Summary of Electrochemical Cell Components

Component	Function
Electrodes	Sites of oxidation/reduction
Electrolyte	Contains ions that carry current
Salt bridge	Maintains electrical neutrality
External wire	Allows flow of electrons

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📘 Chapter 3: Electrochemistry

🧾 Page 3 – Summary Notes (NCERT)

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⚡ Cell Potential and Gibbs Free Energy

• The electrical work done by the electrochemical cell can be related to Gibbs free energy (ΔG).

• Relation Between ΔG and Cell Potential (E_cell):

  $$\Delta G = -nFE_{\text{cell}}$$

  Where:

  • ΔG = Gibbs free energy change

  • n = number of electrons transferred in the reaction

  • F = Faraday’s constant (96,500 C mol⁻¹)

  • E_cell = cell potential (volts)

• Significance:

  • If E_cell > 0, then ΔG < 0 → reaction is spontaneous.

  • If E_cell < 0, then ΔG > 0 → reaction is non-spontaneous.

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🔍 Reversible and Irreversible Cells

• Reversible Electrochemical Cell:

  • Capable of doing electrical work reversibly.

  • The direction of current and reaction can be reversed by applying a slightly higher or lower potential.

• Irreversible Cell:

  • Cannot reverse the reaction just by reversing potential.

  • Irreversibility arises due to:

    • Irreversible chemical changes

    • Side reactions

    • Large overpotentials

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🧠 Example: Daniell Cell

• A classic example of a galvanic cell.

• Uses zinc and copper electrodes in respective sulfate solutions.

• Reaction:

  $$\text{Zn (s)} + \text{Cu}^{2+} (aq) \rightarrow \text{Zn}^{2+} (aq) + \text{Cu (s)}$$
  

• Electrons flow from zinc (anode) to copper (cathode).

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📏 Standard Electrode Potential

• Every half-cell has a standard electrode potential (E°), measured under standard conditions:

  • 1 M concentration

  • 1 atm pressure

  • 25°C temperature

• Measured relative to the Standard Hydrogen Electrode (SHE).

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📘 Chapter 3: Electrochemistry

🧾 Page 4 – Summary Notes (NCERT)

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🧪 Standard Hydrogen Electrode (SHE)

• Acts as a reference electrode with an assigned standard electrode potential (E°) of 0.00 V.

• Construction:

  • A platinum electrode is coated with platinum black.

  • It is immersed in a solution containing 1 M H⁺ ions.

  • Hydrogen gas at 1 atm pressure is bubbled through the solution.

• Half-cell reaction:

  $$\text{H}^{+}(aq) + e⁻ \rightleftharpoons \dfrac{1}{2} \text{H}_2(g)$$

• It can act as either:

  • Anode (oxidation of H₂ to H⁺)

  • Cathode (reduction of H⁺ to H₂)
    depending on the other half-cell it is connected to.

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⚙️ Measuring Electrode Potential of a Half Cell

• Electrode potential cannot be measured in isolation.

• It is always measured relative to another electrode, usually SHE.

💡 Example:

• To measure the standard electrode potential of a Zn²⁺/Zn electrode:

  • Combine it with the SHE.

  • Observe the direction of electron flow:

    • If electrons flow from Zn to SHE, Zn is more reactive (stronger reducing agent).

  • The measured cell potential is taken as the standard reduction potential for the Zn²⁺/Zn couple.

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🔋 Representation of Electrochemical Cell

• General notation:

  $$\text{Anode (oxidation)} \ | \ \text{Electrolyte} \ || \ \text{Electrolyte} \ | \ \text{Cathode (reduction)}$$

• Example:
  Zn | Zn²⁺ (1 M) || Cu²⁺ (1 M) | Cu

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📘 Chapter 3: Electrochemistry

🧾 Page 5 – Summary Notes (NCERT)

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⚙️ Nernst Equation

• The Nernst equation relates the electrode potential to the concentration of ions in a solution and temperature.

• General form of the Nernst equation:

  $$E = E^\circ - \dfrac{0.0591}{n} \log \dfrac{[\text{Red}]_{\text{cathode}}}{[\text{Ox}]_{\text{anode}}}$$

  Where:

  • E = Electrode potential at non-standard conditions

  • E° = Standard electrode potential

  • n = Number of electrons involved in the half-reaction

  • [Red] = Concentration of the reduced species

  • [Ox] = Concentration of the oxidized species

• When the concentrations of the reactants and products are 1 M, the Nernst equation simplifies to:

  $$E = E^\circ$$

• The Nernst equation is useful to calculate the electrode potential under non-standard conditions, especially when concentrations of ions vary.

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⚡ Applications of Nernst Equation

• Cell Potential Calculation:

  • Nernst equation helps in calculating the potential of electrochemical cells under varying concentrations of reactants and products.

• Equilibrium Constant (K):

  • The Nernst equation also provides a relationship between cell potential (E) and the equilibrium constant (K) of a reaction:

  $$E^\circ = \dfrac{0.0591}{n} \log K$$

  • A positive E° indicates a favorable reaction with a large K (products favored).

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🔥 Importance of Nernst Equation

• It allows calculation of cell potential for reactions involving different concentrations.

• Helps predict the spontaneity of reactions.

• Determines the equilibrium position of redox reactions.

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💡 Example of Nernst Equation in Action

• For a Zn/Cu cell:

  • Standard electrode potentials for zinc and copper are given.

  • Using the Nernst equation, the potential of the cell can be determined when the concentrations of zinc and copper ions are not 1 M.