Grade 11
  1. Basic concepts of chemistry  1.1. Importance of Chemistry  1.2. Nature and scope of chemistry  1.3. laws of chemical combination  1.3.1. Law of conservation of mass  1.3.2. Law of definite proportions  1.3.3. Law of multiple proportions  1.3.4. Law of gaseous volume  1.3.5. Avogadro's Law  1.4. Dalton's atomic theory  1.5. Mole concept and molar mass  1.6. Percent composition  1.7. Empirical and molecular formula  1.8. Stoichiometry and Stoichiometric Calculations  1.9. Limiting reagent concept
  2. Structure of the atom  2.1. Discovery of the electron, proton, and neutron  2.2. Atomic Model  2.2.1. Thomson's model  2.2.2. Rutherford's model  2.2.3. Bohr's model  2.3. Quantum mechanical model of the atom  2.4. Dual nature of matter and radiation  2.5. Heisenberg's uncertainty principle  2.6. Quantum numbers  2.7. Atomic orbitals and their shapes  2.8. Electronic configuration of elements  2.9. Aufbau principle  2.10. Pauli's exclusion principle  2.11. Hund's maximum multiplicity rule  2.12. de Broglie's hypothesis  2.13. Photoelectric effect
  3. Classification of elements and periodicity in properties  3.1. Historical development of the periodic table  3.2. Modern Periodic Law and Periodic Table  3.3. Periodic trends in properties  3.3.1. Atomic and Ionic Radius  3.3.2. Ionization enthalpy  3.3.3. Electron gain enthalpy  3.3.4. Electronegativity  3.3.5. Valency  3.3.6. Metallic and Non-Metallic Properties  3.3.7. Oxidation states  3.4. Unusual Properties of the First Elements
  4. Chemical Bonding and Molecular Structure  4.1. Kossel–Lewis approach to chemical bonding  4.2. Ionic or electrovalent bond  4.3. Covalent bond and its features  4.4. Bond parameters  4.5. VSEPR Theory  4.6. Valence bond theory  4.7. Hybridization and its types  4.8. Molecular orbital theory  4.8.1. Bond order and stability  4.9. Hydrogen bonding
  5. States of matter  5.1. Intermolecular forces  5.2. Gas Laws  5.2.1. Boyle's law  5.2.2. Charles's Law  5.2.3. Avogadro's Law  5.2.4. Dalton's law of partial pressure  5.2.5. Graham's law of spread  5.3. Ideal Gas Equation and Applications  5.4. Real gases and deviations from ideal behaviour  5.5. Kinetic molecular theory of gases  5.6. Liquefaction of gases  5.7. Properties of Liquids  5.8. Surface tension and viscosity
  6. Thermodynamics  6.1. System and environment  6.2. Types of systems in thermodynamics  6.3. Types of procedures  6.4. First law of thermodynamics  6.5. Internal energy and enthalpy  6.6. Heat capacity and specific heat  6.7. Hess's law of constant heat summation  6.8. Bond dissociation enthalpy  6.9. Enthalpy of formation and combustion  6.10. Enthalpy of solution and neutralization  6.11. Spontaneity and the second law of thermodynamics  6.12. Gibbs free energy and its applications
  7. Balance  7.1. The concept of balance  7.2. Equilibrium constant and its applications  7.3. Le Chatelier's principle  7.4. Ionic equilibrium in solutions  7.5. Acid and Base Theory  7.5.1. Arrhenius concept  7.5.2. Bronsted-Lowry concept  7.5.3. Lewis concept  7.6. pH Scale and pOH  7.7. Common ion effect  7.8. Hydrolysis of salts  7.9. Buffer solutions and their mechanism of action  7.10. Solubility equilibrium of sparingly soluble salts
  8. Redox reactions  8.1. Oxidation and reduction concepts  8.2. Oxidation number and its applications  8.3. Redox reactions in terms of electron transfer  8.4. Balancing redox reactions (oxidation number and ion-electron methods)  8.5. Types of redox reactions  8.6. Electrochemical Cells and Redox Reactions
  9. Hydrogen  9.1. Position of Hydrogen in the Periodic Table  9.2. Isotopes of hydrogen  9.3. Preparation of Hydrogen  9.4. Properties of Hydrogen  9.5. Hydrides and their classification  9.6. Water and heavy water  9.7. Hydrogen peroxide and its properties  9.8. Uses of Hydrogen and Its Compounds
  10. S-block elements (alkali and alkaline earth metals)  10.1. Group 1 Elements - Properties and Trends  10.2. Group 2 Elements - Properties and Trends  10.3. Unusual properties of lithium and beryllium  10.4. Important compounds of sodium and their uses  10.5. Important Compounds of Magnesium and Calcium
  11. Some p-block elements  11.1. General Introduction to p-Block Elements  11.2. Group 13 Elements  11.2.1. Boron and its compounds  11.3. Group 14 Elements  11.3.1. Carbon and its allotropes  11.3.2. Important Compounds of Carbon and Silicon
  12. Organic Chemistry - Some Basic Principles and Techniques  12.1. Classification and Nomenclature of Organic Compounds  12.2. Structural representation of organic compounds  12.3. Classification of organic reactions  12.4. Electronic Effects in Organic Chemistry  12.4.1. Inductive effect  12.4.2. Resonance effect  12.4.3. Hyperconjugation  12.5. Reaction mechanism and intermediate species  12.6. Purification and qualitative analysis of organic compounds
  13. Hydrocarbons  13.1. Hydrocarbons  13.1.1. Preparation and Properties of Alkenes  13.2. alkene  13.2.1. Preparation and Properties of Alkenes  13.2.2. Electrophilic addition reactions  13.3. Alkynes  13.3.1. Preparation and Properties of Alkynes  13.4. Aromatic hydrocarbons  13.4.1. Structure and properties of benzene  13.4.2. Electrophilic substitution reactions of benzene
  14. Environmental Chemistry  14.1. Environmental Pollution  14.2. Atmospheric pollution and the greenhouse effect  14.3. Water pollution and treatment methods  14.4. Soil pollution and its effects  14.5. Industrial waste and its treatment  14.6. Strategies for pollution control

Grade 11 → Redox reactions ↓


Electrochemical Cells and Redox Reactions


Redox reactions, also called oxidation-reduction reactions, involve the transfer of electrons between two substances. In these reactions, one substance undergoes oxidation while the other undergoes reduction. Understanding these reactions is important in the study of electrochemical cells, which are systems that can convert chemical energy into electrical energy and vice versa.

What are redox reactions?

Redox reactions are characterized by the transfer of electrons from one chemical species to another. Here, the term oxidation means the loss of electrons, while reduction means the gain of electrons. These reactions are always coupled because if one species is oxidized (loses electrons), the other must be reduced (gain electrons).

Example of a redox reaction

Consider the reaction between zinc and copper ions:

        Zn(s) + Cu 2+ (aq) → Zn 2+ (aq) + Cu(s)

In this reaction, zinc loses two electrons and is oxidized to Zn 2+. The copper ions gain those two electrons and are reduced to copper metal.

Oxidation half-reaction

        Zn(s) → Zn2 + (aq) + 2e-

The zinc metal gets oxidised, that is, electrons are removed from it.

Reduction half reaction

        2Cu2 + (aq) + 2e− → Cu(s)

The copper ions are reduced, that is, they gain electrons.

Identifying oxidizing and reducing agents

In a redox reaction, the substance that is oxidized is the reducing agent because it donates electrons. In contrast, the substance that is reduced is the oxidizing agent because it accepts electrons.

In the above example:

  • Zn is a reducing agent.
  • Cu 2+ is the oxidizing agent.

Electrochemical cells

Electrochemical cells are devices that are able to generate electrical energy from chemical reactions or facilitate chemical reactions through the introduction of electrical energy. These cells are divided into two main types:

  • Galvanic (or voltaic) cells
  • Electrolytic cell

Galvanic cells

Galvanic cells derive their power from spontaneous redox reactions. Let's take a closer look at how a typical galvanic cell operates using the zinc-copper example discussed earlier.

Daniell cell

A classic example of a galvanic cell is the Daniell cell, which consists of two half-cells connected by a salt bridge. One half-cell contains a strip of zinc placed in a solution of zinc sulfate, while the other half-cell contains a strip of copper placed in a solution of copper sulfate.

        Zn(s) | ZnSO 4 (aq) || CuSO4 (aq) Cu(s)

Here, the zinc electrode acts as the anode where the oxidation takes place:

        Zn(s) → Zn2 + (aq) + 2e-

The copper electrode acts as the cathode where reduction takes place:

        2Cu2 + (aq) + 2e− → Cu(s)
Zn|Zn 2+ Cu2 + salt bridge

The electrons flow from the zinc electrode to the copper electrode through an external circuit, producing an electric current. The salt bridge (usually a tube containing a gel with ions such as NaNO 3 or KCl) allows ions to move between the two solutions to maintain charge balance.

Electrolytic cell

Unlike galvanic cells, electrolytic cells require external electrical energy to drive non-spontaneous chemical reactions. This type of cell is used in processes such as electrolysis, where compounds are broken down into their constituent elements.

Example: Electrolysis of water

Electrolysis of water is a process in which electric current is passed through water to decompose it into oxygen and hydrogen gases.

        2H 2 O(l) → 2H 2 (g) + O 2 (g)

The setup for this process consists of two electrodes immersed in water. On applying an external voltage, the following reactions take place:

  • Oxidation of water at the anode (positive electrode) produces oxygen gas and hydrogen ions:
    •             2H 2 O(l) → O 2 (g) + 4H + (aq) + 4e -
  • At the cathode (negative electrode), reduction involves hydrogen ions gaining electrons to form hydrogen gas:
    •             4H + (aq) + 4e - → 2H 2 (g)

The overall reaction represents the decomposition of water into gaseous hydrogen and oxygen.

Applications of electrochemical cells

Electrochemical cells have numerous applications. Some of the most important areas are as follows:

Batteries

Batteries are essentially galvanic cells that store chemical energy to convert it into electrical energy. Common types of batteries are:

  • Alkaline batteries
  • Lead-acid batteries (used in vehicles)
  • Lithium-ion batteries (used in electronics)

Each type of battery works on the principles of redox reaction to provide electricity.

Electroplating

Electroplating uses an electrolytic cell to deposit a thin layer of metal on the surface of an object. For example, chromium can be plated on iron to prevent rusting, or gold can be plated on jewellery for aesthetic purposes.

Industrial processes

Electrochemical cells are important in a variety of industrial processes such as:

  • Production of chlorine gas and sodium hydroxide through electrolysis of brine.
  • Manufacture of aluminium from bauxite via the Hall–Héroult process.

Conclusion

The study of electrochemical cells and redox reactions provides insight into how chemical changes can generate or be driven by electrical energy. Understanding these concepts is fundamental to exploring various applications in everyday life and industrial practices.


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Electrochemical Cells and Redox Reactions

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