Grade 10
  1. Matter and its properties  1.1. States of matter  1.2. Physical and chemical properties of matter  1.3. Classification of substances (elements, compounds, mixtures)  1.4. Changes in Matter (Physical and Chemical Changes)  1.5. Law of conservation of mass and law of definite proportions
  2. Atomic Structure  2.1. Historical development of the atomic model  2.2. Subatomic particles (protons, neutrons, electrons)  2.3. Atomic number, mass number and isotopes  2.4. Electronic Configuration and Energy Levels  2.5. Valency, ion formation and oxidation state
  3. Periodic table  3.1. History and Development of the Periodic Table  3.2. Modern Periodic Law and Periodic Trends  3.3. Groups and Periods in the Periodic Table  3.4. Trends in the Periodic Table  3.5. Metals, non-metals, metalloids and their properties  3.6. Noble gases and their chemical inertness
  4. Chemical bond  4.1. Formation of Chemical Bonds (Octet Rule)  4.2. Ionic Bonding and Properties of Ionic Compounds  4.3. Covalent Bond and Properties of Covalent Compounds  4.4. Metallic bond and its properties  4.5. Molecular geometry and VSEPR theory  4.6. Polarity and dipole moment of molecules  4.7. Intermolecular forces
  5. Chemical Reactions and Equations  5.1. Types of Chemical Reactions  5.2. Writing and balancing chemical equations  5.3. Law of conservation of mass in chemical reactions  5.4. Factors Affecting Chemical Reactions  5.5. Catalysts and their role in chemical reactions  5.6. Exothermic and Endothermic Reactions
  6. Stoichiometry and Chemical Calculations  6.1. Mole concept and Avogadro number  6.2. Molar Mass and Gram-Mole Conversions  6.3. Empirical and molecular formula  6.4. Stoichiometric calculations and chemical yield  6.5. Limiting and Excess Reactants
  7. Acids, Bases and Salts  7.1. Properties and Definitions of Acids and Bases (Arrhenius, Bronsted-Lowry, Lewis)  7.2. pH Scale, pOH, and Indicators  7.3. Neutralization reactions and salt formation  7.4. Strength of Acids and Bases (Strong vs. Weak Acids and Bases)  7.5. Common Acids and Bases in Daily Life  7.6. Salts, their solubility and applications
  8. Metals and Nonmetals  8.1. Physical and chemical properties of metals  8.2. Physical and Chemical Properties of Non-Metals  8.3. Reactivity Series of Metals and Displacement Reactions  8.4. Extraction of metals from ores  8.5. Corrosion, its causes and prevention  8.6. Alloys, their composition and uses
  9. Carbon and its compounds  9.1. Allotropes of Carbon  9.2. Hydrocarbons and their classification (alkanes, alkenes, alkynes)  9.3. Functional groups in organic compounds  9.4. Nomenclature of organic compounds (IUPAC system)  9.5. Isomerism in organic chemistry (structural and geometrical)  9.6. Important organic reactions (combustion, addition, substitution, polymerization)  9.7. Applications of organic compounds in industry and daily life
  10. Environmental Chemistry  10.1. Air Pollution (Sources, Effects and Control)  10.2. Water Pollution (Causes, Effects and Remedies)  10.3. Greenhouse effect and global warming  10.4. Ozone layer depletion and its effects  10.5. Green Chemistry and Its Applications  10.6. sustainable chemistry practice
  11. Thermochemistry  11.1. Heat and Energy Changes in Chemical Reactions  11.2. Exothermic and Endothermic Reactions  11.3. Enthalpy, enthalpy change, and heat of reaction  11.4. Specific heat capacity and calorimetry  11.5. Energy Changes in Combustion Reactions
  12. Electrochemistry  12.1. Electrolytes and non-electrolytes  12.2. Electrolysis and its applications (electroplating, extraction of metals)  12.3. Redox reactions (oxidation and reduction)  12.4. Galvanic cell and electrochemical series  12.5. Corrosion and electrochemical prevention methods
  13. Chemical kinetics and equilibrium  13.1. Rate of chemical reactions and its measurement  13.2. Factors affecting the reaction rate (catalyst, temperature, concentration)  13.3. Reversible and irreversible reactions  13.4. Dynamic equilibrium and equilibrium constant  13.5. Le Chatelier's principle and its applications
  14. Gases and Gas Laws  14.1. Properties of gases and kinetic molecular theory  14.2. Boyle's Law (Pressure-Volume Relation)  14.3. Charles's Law  14.4. Gay-Lussac's Law  14.5. Combined Gas Law and Ideal Gas Law  14.6. Real Gases vs Ideal Gases
  15. Nuclear Chemistry  15.1. Introduction to Radioactivity and Nuclear Reactions  15.2. Types of radioactive decay (alpha, beta, gamma)  15.3. Half-life and applications of radioactive isotopes  15.4. Nuclear fission and fusion reactions  15.5. Nuclear power generation and its environmental impact

Grade 10 → Electrochemistry ↓


Redox reactions (oxidation and reduction)


In chemistry, understanding reactions is important to know how substances interact and are transformed. One of the most fundamental types of chemical reactions is a redox reaction, which is short for reduction-oxidation reaction. These are processes that involve the transfer of electrons between two substances. Let's discuss this topic in detail, see what redox reactions are, what their components are, how they work, and their relevance in real-world applications.

What is a redox reaction?

Redox reactions are chemical reactions where the oxidation state of molecules, atoms or ions changes by gaining or losing electrons. The term 'redox' comes from two concepts: reduction and oxidation.

Key concepts

  • Oxidation: This involves the loss of electrons. When a substance loses electrons, it becomes oxidized.
  • Reduction: It involves the gain of electrons. When a substance gains electrons, it is reduced.

Example: Formation of water from the reaction between hydrogen and oxygen.

2H 2 + O 2 → 2H 2 O

In the example above, hydrogen is oxidized (loses electrons), and oxygen is reduced (gains electrons).

Visual example of a redox reaction: electron transfer Hydrogen ( H2 ) - oxidation Oxygen ( O2 ) - Reduction Electron Transfer

Rules for determining oxidation and reduction

To determine what is oxidized and what is reduced in a reaction, we look at oxidation numbers. These numbers help keep track of the electrons in atoms. Here are some basic rules:

  • The oxidation number of an element in its natural state (e.g., O 2 , H 2 ) is zero.
  • The oxidation number of a monatomic ion is the same as its charge (for example, Na + is +1).
  • The oxidation number of oxygen is usually -2, except in peroxides such as H 2 O 2.
  • The oxidation number of hydrogen is normally +1, except when it is bonded to metals in hydrides (e.g., LiH).
  • The sum of the oxidation numbers in a neutral compound is zero; in a polyatomic ion, it is equal to the charge of the ion.
Example of assigning oxidation numbers: 
MnO 4 - :
    - O = -2 (4 oxygen = -8 total)
    - Total charge is -1.
    - Mn = +7 (total oxidation = -1)

Identifying redox reactions

Not all reactions are redox reactions. To determine whether a reaction is a redox reaction, check to see if there is a change in the oxidation numbers. Consider the reaction between zinc and copper (II) sulfate:

4Zn + CuSO4 → ZnSO4 + Cu

In this reaction, zinc changes from 0 in Zn to +2 in ZnSO4 (oxidation), and copper changes from +2 in CuSO4 to 0 in Cu (reduction).

Visual example of redox process Zn(0) Zn 2+ (+2) Cube 2+ (+2) Cube (0)

Balancing redox reactions

Balancing redox reactions ensures conservation of mass and charge. Let's use the ion-electron method, which is especially useful in acidic or alkaline solutions:

  1. Divide the reaction into oxidation and reduction half-reactions.
  2. Balance each half-reaction for mass and charge.
  3. Combine the half-reactions, making sure that the electrons cancel out.

Example: Balancing a redox reaction in an acidic solution

Consider balancing the following reaction:

MnO 4 - + Fe 2+ → Mn 2+ + Fe 3+
  1. Write the half-reactions: 
    Oxidation: Fe 2+ → Fe 3+ + e -
Reduction: MnO 4 - + 8H + + 5e - → Mn 2+ + 4H 2 O
  2. Balance the electrons: 
    Multiply the oxidation half-reaction by 5:
5Fe2 + → 5Fe3 + + 5e-

Mix:
5Fe 2+ + MnO 4 - + 8H + → 5Fe 3+ + Mn 2+ + 4H 2 O

Applications of redox reactions

Redox reactions are not just theoretical concepts but also have real applications in various fields:

Batteries

Redox reactions are the core of battery function. In a battery, redox reactions occur in an electrochemical cell, where oxidation occurs at the anode, and reduction occurs at the cathode, producing electrical energy.

Example: Lead-acid battery

Pb + PbO 2 + 2H 2 SO 4 → 2PbSO 4 + 2H 2 O

Corrosion

Rust is an example of an undesirable redox reaction. Iron reacts with oxygen and moisture present in the atmosphere, causing rust.

4Fe + 3O 2 + 6H 2 O → 4Fe(OH) 3

Metabolism and respiration

Biological systems depend on redox reactions for energy. Cellular respiration is a redox reaction in which glucose is oxidized, yielding energy.

C 6 H 12 O 6 + 6O 2 → 6CO 2 + 6H 2 O + energy

Environmental chemistry

Redox reactions play an important role in purifying pollutants and maintaining environmental balance. For example, water treatment involves redox processes that remove harmful pollutants.

Conclusion

Understanding redox reactions is essential in chemistry because they are fundamental to processes in nature, industry, and technology. Identifying how electrons are transferred enables us to better understand energy conversion, chemical synthesis, and biological systems. From the corrosion of metals to battery operation, redox reactions are indispensable for a variety of applications that are vital to modern life.

Mastering the rules and concepts of oxidation and reduction, balancing redox reactions, and applying these principles to real-world situations is vital for any chemistry student. As you progress in your chemistry education, these fundamental concepts will reappear and expand upon, making their importance in the study of science even stronger.


Grade 10 → 12.3


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Redox reactions (oxidation and reduction)

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