Electrochemistry deals with the relationship between chemical reactions and electrical energy. It explains how chemical energy is converted into electrical energy and vice versa.

👉 Core Idea: Redox reactions can produce electricity or be driven by electricity.

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1. Electrochemical Cells

Definition

An electrochemical cell is a device that converts chemical energy into electrical energy.

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Types of Cells

(A) Galvanic Cell (Voltaic Cell)

• Produces electricity
• Based on spontaneous reaction

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(B) Electrolytic Cell

• Uses electricity
• Non-spontaneous reaction

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Example: Daniell Cell

Representation

Zn | Zn²⁺ || Cu²⁺ | Cu

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Diagram (Conceptual)

Zn → Zn²⁺ + 2e⁻      (Anode)
Cu²⁺ + 2e⁻ → Cu      (Cathode)Electron flow: Zn → Cu
Salt bridge connects solutions

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Concept Clarity

👉 WHY salt bridge is used?
To maintain electrical neutrality.

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2. Electrode Potential

Definition

The tendency of an electrode to lose or gain electrons.

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Standard Electrode Potential (E°)

Measured under standard conditions:

• 1 M concentration
• 1 atm pressure
• 298 K temperature

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Standard Hydrogen Electrode (SHE)

Reference electrode with E° = 0 V

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3. EMF of a Cell

Definition

The potential difference between two electrodes.

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Formula

E°cell = E°cathode − E°anode

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Concept Clarity

👉 Positive EMF → spontaneous reaction

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4. Nernst Equation (Very Important)

Formula

E = E° − (0.0591/n) log Q

Where:

• n = number of electrons
• Q = reaction quotient

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Application

• Calculate EMF under non-standard conditions
• Determine equilibrium constant

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5. Conductance

Definition

Ability of a solution to conduct electricity.

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Resistance

R = V / I

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Conductance

G = 1 / R

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Specific Conductance (κ)

Conductance of 1 cm³ solution

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Molar Conductance

Λₘ = κ × 1000 / C

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Concept Clarity

👉 WHY conductance increases on dilution?
Because ions move more freely.

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6. Variation of Conductance

Strong Electrolytes

• Fully ionized
• Slight increase on dilution

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Weak Electrolytes

• Partially ionized
• Sharp increase on dilution

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7. Kohlrausch’s Law

Statement

At infinite dilution, each ion contributes independently to conductance.

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Formula

Λₘ° = λ⁺ + λ⁻

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8. Electrolysis

Definition

Chemical decomposition using electric current.

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Faraday’s Laws of Electrolysis

First Law

Mass ∝ quantity of electricity

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Second Law

Mass ∝ equivalent weight

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Formula

m = (ZIt)

Where:

• Z = electrochemical equivalent

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9. Batteries

Primary Cells

• Non-rechargeable
  Example: Dry cell

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Secondary Cells

• Rechargeable
  Example: Lead-acid battery

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Fuel Cells

Convert chemical energy directly into electricity

Example: H₂–O₂ fuel cell

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10. Corrosion

Definition

Slow deterioration of metals due to chemical reactions.

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Example

Rusting of iron

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Prevention

• Painting
• Galvanization
• Alloying

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11. Important Numericals

Numerical 1

Find EMF if E°cathode = 0.34 V, E°anode = -0.76 V

E°cell = 1.10 V

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Numerical 2

Calculate EMF using Nernst equation

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Numerical 3

Find molar conductance

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Numerical 4

Find mass deposited using Faraday law

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12. Important Formula Sheet

• E°cell = E°cathode − E°anode
• Nernst equation
• Λₘ = κ × 1000 / C
• m = ZIt

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13. Concept Clarity (Very Important)

👉 WHY electrons flow from anode to cathode?
Because anode undergoes oxidation.

👉 WHY EMF decreases over time?
Because concentration changes.

👉 WHY strong electrolytes conduct better?
Because they have more ions.

👉 WHY corrosion occurs?
Because metals react with environment.

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14. Common Mistakes

• Confusing anode and cathode
• Wrong sign in EMF formula
• Errors in Nernst equation

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Conclusion

Electrochemistry connects chemistry with electricity and has vast applications in batteries, corrosion prevention, and industrial processes.

👉 Focus on:

• Nernst equation
• EMF calculations
• Conductance concepts