Unit 3

General and Physical Chemistry

Class 12 Chemistry

Electrochemistry

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Curriculum alignment note: The legacy Nepal eNotes source page displays “Unit 3” for this post. In the Grade 12 Chemistry curriculum, Electrochemistry is listed as Unit 5 under General and Physical Chemistry. The typed notes below follow that curriculum scope: electrode potential, reference electrodes, electrochemical series, voltaic cells, cell potential, free-energy relation, and commercial cells/fuel cells.

1. Introduction to Electrochemistry

Definition Electrochemistry is the branch of chemistry that studies the relationship between chemical change and electrical energy, especially the interconversion of chemical energy and electrical energy through redox processes.

In a spontaneous electrochemical cell, a redox reaction produces electrical energy. Oxidation and reduction occur at separate electrodes so that electrons are forced to travel through an external circuit. This controlled electron flow is electric current.

Oxidation

Loss of electrons. It occurs at the anode.

Zn(s) → Zn²⁺(aq) + 2e⁻

Reduction

Gain of electrons. It occurs at the cathode.

Cu²⁺(aq) + 2e⁻ → Cu(s)

Important Note For a voltaic (galvanic) cell, the spontaneous reaction sends electrons from anode to cathode through the wire. The salt bridge completes the internal ionic circuit and helps maintain electrical neutrality.

2. Electrode Potential

When a metal is placed in a solution containing its ions, two opposing tendencies are possible:

metal atoms may enter solution as positive ions, leaving electrons on the metal;
metal ions in solution may gain electrons and deposit on the metal surface.

The charge separation that develops at the metal–solution interface produces a potential difference called electrode potential.

Definition Electrode potential is the potential difference developed between an electrode and the solution of its ions when equilibrium is established at their interface.

Diagram 1: Origin of electrode potential at a metal–ion interface

Remember The absolute potential of a single electrode is not measured directly. It is measured relative to a reference electrode, commonly the Standard Hydrogen Electrode.

3. Standard Electrode Potential (E°)

When an electrode is measured under standard conditions, the resulting electrode potential is called the standard electrode potential.

Standard conditions A solution is taken at unit activity (commonly represented in school calculations as approximately 1 mol L⁻¹), gases at standard pressure, pure solids/liquids in their standard states, usually at 298 K.

Tabulated electrode potentials are normally written as standard reduction potentials. For example:

Cu²⁺(aq) + 2e⁻ → Cu(s) E° = +0.34 V Zn²⁺(aq) + 2e⁻ → Zn(s) E° = −0.76 V

Interpretation A more positive standard reduction potential means a greater tendency for the species on the left side of the reduction half-reaction to gain electrons.

4. Standard Hydrogen Electrode (SHE)

The Standard Hydrogen Electrode is the primary reference electrode. Its standard electrode potential is assigned exactly:

Reference value E°(H⁺/H₂) = 0.00 V 2H⁺(aq) + 2e⁻ ⇌ H₂(g)

Construction

An inert platinum electrode coated with platinum black is used.
Hydrogen gas is passed over the platinum surface.
The electrode contacts a solution containing H⁺ ions under standard conditions.
It is connected to the other half-cell through an external circuit and salt bridge.

Diagram 2: Standard Hydrogen Electrode (SHE)

Exam Important Be able to draw and label the SHE, write its half-reaction, state its standard potential, and explain why platinum is used: it is inert, electrically conducting, and provides a surface for the H₂/H⁺ equilibrium.

5. Calomel Electrode

The calomel electrode is a practical reference electrode based on mercury, mercury(I) chloride (calomel), and potassium chloride solution.

Electrode representation Hg(l) | Hg₂Cl₂(s) | KCl(aq) Hg₂Cl₂(s) + 2e⁻ ⇌ 2Hg(l) + 2Cl⁻(aq)

Its potential depends on the chloride-ion concentration. Therefore, the KCl concentration must be specified when precise values are required.

Diagram 3: Simplified calomel electrode

Feature	Standard Hydrogen Electrode	Calomel Electrode
Reference system	H⁺/H₂	Hg₂Cl₂/Hg with Cl⁻
Assigned/reference role	Primary reference; E° = 0.00 V	Secondary practical reference
Main electrode material	Platinized platinum	Mercury in contact with calomel
Important condition	Standard H⁺ and H₂ conditions	KCl concentration must be known

6. Electrochemical Series and Its Applications

The electrochemical series is an arrangement of redox couples according to their standard reduction potentials.

Reduction half-reaction	Typical E° (V)	Meaning
Ag⁺ + e⁻ → Ag	+0.80	Ag⁺ has strong tendency to be reduced relative to H⁺.
Cu²⁺ + 2e⁻ → Cu	+0.34	Cu²⁺ is reduced more readily than H⁺ under standard conditions.
2H⁺ + 2e⁻ → H₂	0.00	Reference point.
Zn²⁺ + 2e⁻ → Zn	−0.76	Zn metal is a relatively strong reducing agent.

Major Applications

Identify cathode and anode: the half-cell with the more positive reduction potential acts as the cathode in a spontaneous cell.
Predict electron flow: electrons move externally from the anode to the cathode.
Calculate standard cell potential: use the difference between cathode and anode reduction potentials.
Predict feasibility: a positive E°cell corresponds to a spontaneous cell reaction under standard conditions.
Compare oxidizing strength: a more positive reduction potential generally means a stronger oxidizing agent.
Compare reducing strength: the reduced form of a very negative couple is generally a stronger reducing agent.

Common mistake Do not multiply an electrode potential by a stoichiometric coefficient. E° is an intensive quantity. Balance electrons in the half-reactions, but use the tabulated E° value unchanged.

7. Voltaic (Galvanic) Cell

Definition A voltaic or galvanic cell is an electrochemical cell in which a spontaneous redox reaction converts chemical energy into electrical energy.

Part	Function in a voltaic cell
Anode	Oxidation occurs; it is the negative electrode in a voltaic cell.
Cathode	Reduction occurs; it is the positive electrode in a voltaic cell.
External wire	Carries electrons from anode to cathode.
Salt bridge	Allows ion migration, maintains electrical neutrality, and completes the circuit.
Voltmeter/load	Measures or uses the cell potential.

7.1 Zn–Cu (Daniell) Cell

The Zn–Cu cell contains a zinc electrode in Zn²⁺ solution and a copper electrode in Cu²⁺ solution.

Anode: Zn(s) → Zn²⁺(aq) + 2e⁻ Cathode: Cu²⁺(aq) + 2e⁻ → Cu(s) Overall: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)

Cell notation Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)

Diagram 4: Zn–Cu voltaic (Daniell) cell

Worked Example 1: Standard EMF of Zn–Cu cell

Given E°(Cu²⁺/Cu) = +0.34 V and E°(Zn²⁺/Zn) = −0.76 V. E°cell = E°cathode − E°anode

Step 1: Cu has the more positive reduction potential, so Cu is the cathode.

Step 2: Zn is the anode.

Step 3: Substitute:

E°cell = (+0.34 V) − (−0.76 V) = +1.10 V

Answer: E°cell = +1.10 V. The positive sign indicates a spontaneous standard cell reaction.

7.2 Ag–Cu Cell

For silver and copper standard half-cells:

Ag⁺ + e⁻ → Ag E° = +0.80 V Cu²⁺ + 2e⁻ → Cu E° = +0.34 V

Ag⁺/Ag has the more positive reduction potential, so silver is the cathode and copper is oxidized at the anode.

Anode: Cu(s) → Cu²⁺(aq) + 2e⁻ Cathode: 2Ag⁺(aq) + 2e⁻ → 2Ag(s) Overall: Cu(s) + 2Ag⁺(aq) → Cu²⁺(aq) + 2Ag(s)

Diagram 5: Ag–Cu voltaic cell

Worked Example 2: E°cell for Ag–Cu cell E°cell = E°cathode − E°anode = 0.80 − 0.34 = +0.46 V

Therefore the standard cell reaction is spontaneous in the direction:

Cu(s) + 2Ag⁺(aq) → Cu²⁺(aq) + 2Ag(s)

8. Cell Potential and Standard Cell Potential

Core formula Ecell = Ecathode − Eanode

When all species are in their standard states:

E°cell = E°cathode − E°anode

Sign of E°cell and Feasibility

E°cell	Meaning under standard conditions
E°cell > 0	Forward cell reaction is spontaneous.
E°cell = 0	System is at equilibrium in the thermodynamic sense.
E°cell < 0	Forward reaction is non-spontaneous; the reverse direction is favored.

Worked Example 3: Predict the direction of reaction Suppose: E°(X²⁺/X) = −0.40 V E°(Y²⁺/Y) = +0.20 V

Y²⁺/Y is more positive, so Y²⁺ is reduced at the cathode. X is oxidized at the anode.

E°cell = +0.20 − (−0.40) = +0.60 V

Because E°cell is positive, X + Y²⁺ → X²⁺ + Y is feasible under standard conditions.

Remember If a table lists reduction potentials, use them directly in E°cathode − E°anode. Do not first change the sign of the anode value and then subtract again; that would reverse the sign twice.

9. Relationship Between Cell Potential and Gibbs Free Energy

The maximum electrical work obtainable from a reversible electrochemical cell is related to Gibbs free-energy change.

Important relation ΔG = −nFE ΔG° = −nFE°

Where:

ΔG = Gibbs free-energy change, usually J mol⁻¹
n = moles of electrons transferred in the balanced cell reaction
F = Faraday constant ≈ 9.6485 × 10⁴ C mol⁻¹
E = cell potential in volts (V = J C⁻¹)

Short Derivation

Electrical work is charge multiplied by potential difference:

electrical work = qE

For a balanced reaction transferring n moles of electrons, the charge is:

q = nF

For a reversible cell, the maximum non-expansion electrical work is related to Gibbs free energy. Using the chemistry sign convention gives:

ΔG = −nFE

Dimensional check (mol e⁻) × (C mol⁻¹ e⁻) × (J C⁻¹) = J So the units are consistent with energy.

Worked Example 4: ΔG° for a Zn–Cu cell

For the Daniell cell, E°cell = 1.10 V and n = 2.

ΔG° = −nFE° = −(2)(9.6485 × 10⁴ C mol⁻¹)(1.10 J C⁻¹) ≈ −2.12 × 10⁵ J mol⁻¹

Answer: ΔG° ≈ −212 kJ mol⁻¹.

The negative ΔG° agrees with the positive E°cell and confirms spontaneity under standard conditions.

E°cell	ΔG°	Spontaneity
Positive	Negative	Forward reaction spontaneous
Zero	Zero	Equilibrium
Negative	Positive	Forward reaction non-spontaneous

10. Commercial Cells, Rechargeable Batteries and Fuel Cells

Electrochemical technology makes it possible to store or continuously generate electrical energy using redox reactions.

10.1 Primary and Secondary Cells

Type	Main idea	Typical feature
Primary cell	Cell reaction is not conveniently reversed during normal use.	Used until reactants are substantially consumed.
Secondary cell	Cell reaction can be reversed by applying an external electrical source.	Rechargeable; lithium-ion batteries are a major modern example.
Fuel cell	Reactants are continuously supplied from outside the cell.	Can operate as long as fuel and oxidant are supplied.

10.2 Lithium-Ion Rechargeable Battery

A lithium-ion cell transfers Li⁺ ions through an electrolyte between two host materials while electrons move through the external circuit. During discharge, the anode releases lithium ions and electrons; during charging, an external power source drives the process in reverse.

Diagram 6: Simplified lithium-ion rechargeable cell

Advantages of rechargeable lithium-ion cells High energy density, low mass for the stored energy, rechargeable operation, and wide use in phones, laptops, power tools, and electric vehicles.

10.3 Hydrogen–Oxygen Fuel Cell

A hydrogen–oxygen fuel cell converts the chemical energy of hydrogen and oxygen directly into electrical energy. The overall product is water.

One convenient acidic representation is:

Anode: 2H₂(g) → 4H⁺(aq) + 4e⁻ Cathode: O₂(g) + 4H⁺(aq) + 4e⁻ → 2H₂O(l) Overall: 2H₂(g) + O₂(g) → 2H₂O(l)

Diagram 7: Hydrogen–oxygen fuel cell

Exam Important A fuel cell differs from an ordinary battery because its reactants are supplied continuously. For the hydrogen–oxygen fuel cell, the overall chemical product is water.

11. Key Formula and Concept Sheet

Concept	Formula / rule	Unit / note
Standard hydrogen electrode	E°(H⁺/H₂) = 0.00	V
Cell potential	Ecell = Ecathode − Eanode	V
Standard cell potential	E°cell = E°cathode − E°anode	V
Free energy relation	ΔG = −nFE	J mol⁻¹ for molar reaction convention
Standard free energy	ΔG° = −nFE°	J mol⁻¹
Faraday constant	F ≈ 9.6485 × 10⁴	C mol⁻¹ e⁻
Spontaneous standard cell	E°cell > 0 and ΔG° < 0	Thermodynamic criterion

12. Common Exam Mistakes

Writing E°anode − E°cathode instead of E°cathode − E°anode.
Multiplying E° by 2 or 3 after multiplying a half-reaction. Electrode potential is not multiplied by stoichiometric coefficients.
Calling the anode positive in a voltaic cell. In a spontaneous voltaic cell, the anode is negative and the cathode is positive.
Drawing electrons through the salt bridge. Electrons travel through the external wire; ions move through the salt bridge.
Forgetting to balance electron transfer before finding n in ΔG = −nFE.
Using the wrong sign of E° and then making a second sign reversal.
Confusing the direction of conventional current with electron flow.
Claiming that a single electrode potential is measured absolutely; it is measured relative to a reference electrode.

13. Quick Comparison: Voltaic Cell vs Reference Electrode

Feature	Voltaic cell	Reference electrode
Purpose	Produces electrical energy from a spontaneous redox reaction	Provides a stable comparison potential
Examples	Zn–Cu, Ag–Cu	SHE, calomel electrode
Measurement	Has a measurable cell EMF	Used as one half-cell to determine another electrode potential
Core idea	Two half-cells operate together	Stable reference half-cell

14. Important Exam Questions

Short-Answer Questions

Define electrochemistry and electrode potential.
What is standard electrode potential?
Why can the absolute potential of a single electrode not be measured directly?
State the standard electrode potential assigned to the Standard Hydrogen Electrode.
Write the representation and electrode reaction of the SHE.
What is a calomel electrode? Write its electrode reaction.
Define electrochemical series.
State four applications of the electrochemical series.
What is a salt bridge? State two functions.
Differentiate between anode and cathode in a voltaic cell.
State the relation between E°cell and ΔG°.
What does a positive E°cell indicate?
What is a fuel cell? How does it differ from an ordinary battery?

Long-Answer Questions

Describe the construction and working of the Standard Hydrogen Electrode with a labelled diagram.
Explain the construction and working principle of a calomel electrode.
Describe a Zn–Cu voltaic cell, including half-reactions, overall reaction, cell notation, electron flow, and salt-bridge function.
Explain an Ag–Cu cell and calculate its standard cell potential from standard electrode potentials.
Explain the electrochemical series and discuss its applications in predicting redox reactions.
Derive and explain the relation ΔG = −nFE.
Describe the principle and working of a hydrogen–oxygen fuel cell with balanced electrode reactions.
Explain the basic working principle and advantages of a lithium-ion rechargeable battery.

Numerical Questions

Given E°(Cu²⁺/Cu) = +0.34 V and E°(Zn²⁺/Zn) = −0.76 V, calculate E°cell for the Zn–Cu cell.
Given E°(Ag⁺/Ag) = +0.80 V and E°(Cu²⁺/Cu) = +0.34 V, calculate E°cell and identify the cathode.
A cell reaction transfers 2 mol of electrons and has E°cell = 0.46 V. Calculate ΔG° using F = 9.6485 × 10⁴ C mol⁻¹.
Two half-cells have reduction potentials −0.25 V and +0.60 V. Determine the spontaneous direction and E°cell.
For a balanced cell reaction with n = 3 and E° = 0.75 V, calculate ΔG° in kJ mol⁻¹.

Diagram Questions

Draw a labelled Standard Hydrogen Electrode.
Draw a labelled calomel electrode.
Draw the Zn–Cu voltaic cell showing anode, cathode, salt bridge, and electron flow.
Draw the Ag–Cu cell and indicate the direction of electron flow.
Draw a simple hydrogen–oxygen fuel-cell diagram.

Exam Strategy In numerical answers, first write the relevant half-cell potentials, identify cathode and anode, write the formula, substitute with signs, include the unit, and finally interpret whether the calculated E°cell is positive or negative.

15. One-Minute Revision

Electrochemistry connects redox reactions with electrical energy.
Oxidation occurs at the anode; reduction occurs at the cathode.
Electrode potential develops at an electrode–solution interface.
Single electrode potentials are measured relative to a reference electrode.
The Standard Hydrogen Electrode has E° = 0.00 V.
The calomel electrode is a practical reference electrode based on Hg/Hg₂Cl₂/KCl.
A more positive standard reduction potential means a greater tendency to be reduced.
In a voltaic cell, electrons flow externally from anode to cathode.
The salt bridge carries ions, maintains neutrality, and completes the circuit.
E°cell = E°cathode − E°anode.
A positive E°cell indicates a spontaneous standard cell reaction.
ΔG° = −nFE°; therefore positive E° means negative ΔG°.
Zn–Cu has E°cell ≈ +1.10 V using common standard potentials.
Ag–Cu has E°cell ≈ +0.46 V using common standard potentials.
Hydrogen–oxygen fuel cells produce water overall; lithium-ion cells are rechargeable.

16. Diagram Practice

Students should practice drawing these diagrams neatly with labels for the NEB examination:

Origin of electrode potential at a metal–ion interface.
Standard Hydrogen Electrode (SHE).
Calomel reference electrode.
Zn–Cu (Daniell) voltaic cell.
Ag–Cu voltaic cell.
Lithium-ion battery schematic.
Hydrogen–oxygen fuel cell.

Source handling: The original Nepal eNotes PDF remains embedded above. The typed section follows the verified NEB/CDC syllabus and is designed as a searchable, responsive study companion. Where the PDF viewer does not expose handwritten page text, the typed section is a syllabus-aligned reconstruction and is not claimed to be a word-for-word transcription.