Grade 12
  1. Solid state  1.1. Classification of solids  1.2. Crystal Lattice and Unit Cell  1.3. Packing Efficiency  1.4. Density of unit cell  1.5. Imperfections in solids (point and line defects)  1.6. Electrical Properties of Solids (Conductors, Insulators and Semiconductors)  1.7. Magnetic Properties of Solids (Diamagnetic, Paramagnetic and Ferromagnetic Materials)
  2. Solution  2.1. Types of Solutions (Homogeneous and Heterogeneous)  2.2. Expressing concentration of solutions (mass %, volume %, molarity, molality, normality, mole fraction)  2.3. Solubility of solids and gases in liquids (Henry's law)  2.4. Raoult's Law and Ideal and Non-Ideal Solutions  2.5. Syndrome properties  2.6. Abnormal molar mass and Van't Hoff factor
  3. Electrochemistry  3.1. Electrochemical cells (galvanic and electrolytic cells)  3.2. Galvanic cell and EMF of the cell  3.3. Standard electrode potential and electrochemical series  3.4. Nernst equation and its applications  3.5. Conductivity and electrolytic cells (specific, molar and equivalent conductivity)  3.6. Kohlrausch's law and its applications  3.7. Batteries (primary and secondary cells)  3.8. Corrosion and its prevention
  4. Chemical kinetics  4.1. Rate of chemical reaction  4.2. Factors affecting the reaction rate (concentration, temperature, catalyst, surface area)  4.3. Order and molecularity of the reaction  4.4. Integrated Rate Equations (Zero, First and Second Order)  4.5. Half-life of a reaction  4.6. Collision theory and activation energy  4.7. Arrhenius equation and its applications
  5. Surface chemistry  5.1. Adsorption and its types (Physicosorption and Chemisorption)  5.2. Catalysis and its types (homogeneous and heterogeneous)  5.3. Colloids and their properties (Tyndall effect, Brownian motion, coagulation)  5.4. Emulsions and gels  5.5. Applications of Colloids
  6. General principles and processes of separation of elements  6.1. Presence of metals and minerals  6.2. Concentration of ores (gravity separation, froth flotation, magnetic separation)  6.3. Extraction of raw metals (roasting, calcination, reduction)  6.4. Thermodynamic and electrochemical principles of metallurgy  6.5. Refining of metals
  7. P-block elements  7.1. Group 15 elements (nitrogen and phosphorus)  7.2. Group 16 Elements (Oxygen, Sulphur and their Compounds)  7.3. Group 17 Elements (Halogens and their Compounds)  7.4. Group 18 Elements  7.5. Properties and Trends in p-Block Elements  7.6. Important compounds of p-block elements (ammonia, phosphine, sulphuric acid, hydrochloric acid)  7.7. Interhalogen compounds and their applications
  8. d-block and f-block elements  8.1. General Properties of Transition Metals  8.2. Properties of Inner Transition Metals (Lanthanoids and Actinoids)  8.3. Lanthanoid and actinoid contractions  8.4. Coordination Chemistry of d-Block Elements
  9. Coordination compounds  9.1. Werner's theory and modern concepts  9.2. Coordination number and type of ligand  9.3. Nomenclature of Coordination Compounds  9.4. Isomerism (geometrical and optical) in coordination compounds  9.5. Bonding in Coordination Compounds (Valence Bond Theory and Crystal Field Theory)  9.6. Stability and applications of coordination compounds in medicine and industry
  10. Haloalkanes and haloarenes  10.1. Classification and nomenclature of haloalkanes and haloarenes  10.2. Physical and Chemical Properties of Haloalkanes and Haloarenes  10.3. Mechanism of Nucleophilic Substitution Reactions (SN1 and SN2)  10.4. Uses and environmental effects (ozone depletion, CFCs)
  11. Alcohol, phenol and ether  11.1. Structure and classification of alcohols, phenols and ethers  11.2. Preparation and properties of alcohols  11.3. Preparation and properties of phenol  11.4. Preparation and properties of ethers  11.5. Chemical Reactions and Uses
  12. Aldehydes, ketones and carboxylic acids  12.1. Structure and nomenclature in aldehydes, ketones and carboxylic acids  12.2. Preparation and properties of aldehydes and ketones  12.3. Chemical reactions in aldehydes, ketones and carboxylic acids (nucleophilic addition, oxidation and reduction)  12.4. Preparation and Properties of Carboxylic Acids  12.5. Chemical Reactions and Uses of Carboxylic Acids
  13. Nitrogen-containing organic compounds  13.1. Amines (classification, structure, basicity)  13.2. Preparation and properties of amines  13.3. Reactions of Amines (Diazotization, Carbylamine Reaction)  13.4. Diazonium salts and their applications in paint industry
  14.1. Carbohydrates (classification and properties)  14.2. Proteins (Structure and Function)  14.3. Enzymes (Mechanism and Function)  14.4. Nucleic acids (DNA and RNA)  14.5. Vitamins and hormones
  15. Polymer  15.1. Classification of Polymers (Addition, Condensation, Copolymers)  15.2. Types of polymerization (free radical, ionic, condensation)  15.3. Important Polymers and Their Uses  15.4. Biodegradable and non-biodegradable polymers
  16. Chemistry in Everyday Life  16.1. Drugs and their classification (analgesics, antipyretics, antibiotics, antacids)  16.2. Chemicals in food and cosmetics (preservatives, artificial sweeteners, antioxidants)  16.3. Cleaning agents and detergents (soaps, synthetic detergents, surfactants)  16.4. Environmental Chemistry (Pollution, Green Chemistry, Sustainable Chemistry)

Grade 12 → Electrochemistry ↓


Electrochemical cells (galvanic and electrolytic cells)


Electrochemistry is an important field of chemistry that deals with the relationship between electrical energy and chemical reactions, particularly oxidation-reduction reactions. At the heart of electrochemistry are electrochemical cells, which can either generate electrical energy from chemical reactions or use electrical energy to drive non-spontaneous reactions.

What are electrochemical cells?

Electrochemical cells are devices that convert chemical energy into electrical energy or vice versa. There are two primary types of electrochemical cells:

  • Galvanic cell (or voltaic cell)
  • Electrolytic cell

Galvanic cells

The galvanic cell is designed to convert chemical energy into electrical energy through spontaneous chemical reactions. It generates electricity when a redox reaction occurs.

Structure of galvanic cell

A basic galvanic cell consists of two different metals immersed in their respective ionic solutions and connected by a wire, allowing electrons to flow. An additional component, the salt bridge, completes the circuit by allowing ion transfer and maintaining charge balance.

        Zn | Zn 2+ (aq) || Cu2+ (aq) cube

In this specific arrangement, zinc is oxidized, losing electrons, while copper is reduced, gaining electrons.

How does a galvanic cell work?

The chemical reactions in a galvanic cell occur as follows:

  • Oxidation reaction (anode): This is where zinc metal (Zn) loses electrons to become zinc ions (Zn → Zn2+ + 2e-).
  • Reduction reaction (cathode): Copper ions in solution gain electrons and form copper metal (Cu2+ + 2e- → Cu).

The flow of electrons from the zinc to the copper electrode through the outer wire produces electrical energy.

Electrode potential

The potential difference between the two half-cells drives the movement of electrons in the circuit. Each metal has an associated standard electrode potential, calculated under standard conditions of 1 M concentration, 25°C temperature, and 1 atm pressure.

By measuring the voltage, we can determine the cell potential (EMF) using the following equation:

        E cell = E cathode - E anode

Electrolytic cell

Unlike galvanic cells, electrolytic cells use electrical energy to drive non-spontaneous chemical reactions. They require an external power source, such as a battery or power supply, to start and sustain the reactions.

Structure of electrolytic cell

An electrolytic cell consists of two electrodes immersed in an electrolyte solution. A power source is connected to the electrodes to make the reaction proceed.

        { Power source } - Anode (Oxidation) | Electrolyte | Cathode (Reduction)

How does an electrolytic cell work?

The main process is electrolysis:

  • Oxidation at the anode: Anions present in the electrolyte lose electrons and move toward the anode. For example, in the electrolysis of molten sodium chloride, chloride ions are oxidized to chlorine gas (2Cl- - 2e- → Cl2).
  • Reduction at the cathode: Cations gain electrons and move toward the cathode. In the same example, sodium ions gain electrons to form sodium metal (Na+ + e- → Na).

Applications of electrolytic cells

Electrolytic cells are widely used in industrial applications, including:

  • Electroplating
  • Electrorefining
  • Production of chemicals such as chlorine and sodium hydroxide

Comparison of galvanic and electrolytic cells

Although both types of cells involve oxidation and reduction reactions, they have clear differences:

Aspect Galvanic cell Electrolytic cell
Energy conversion Chemical to electrical Electrical to chemical
Independence Spontaneous reaction Non-spontaneous reaction
External power Not required Necessary

Visual example of a galvanic cell (simplified)

Zinc Copper

Example calculation for cell potential

To calculate the standard voltage, let us consider a cell made of zinc and copper:

  • E cathode = +0.34 V
  • E anode = -0.76 V

Based on the cell potential equation:

        E cell = E cathode - E anode = 0.34 V - (-0.76 V) = 1.10 V

This positive cell potential indicates a spontaneous reaction, which is typical of a galvanic cell.

Concluding remarks on electrochemical cells

Electrochemical cells play a fundamental role in both chemistry and everyday life. Galvanic cells are essential for the batteries found in countless devices, while electrolytic cells are important in industrial processes and making chemical compounds. Understanding their principles provides the basis for understanding how energy conversion is achieved through chemical reactions.


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Electrochemical cells (galvanic and electrolytic cells)

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