Grade 9
  1. Matter and its nature  1.1. Introduction to matter  1.2. Physical and chemical properties of matter  1.3. States of matter  1.3.1. Solid state  1.3.2. Fluid state  1.3.3. Gaseous state  1.3.4. Plasma and Bose–Einstein condensate  1.4. Changes in the states of matter  1.4.1. Melting and freezing  1.4.2. Evaporation and Condensation  1.4.3. Sublimation and deposition  1.5. Classification of matter  1.6. Elements, Compounds, and Mixtures  1.7. Pure substances and impurities  1.8. Separation Techniques  1.8.1. Filtration  1.8.2. Distillation  1.8.3. Crystallization  1.8.4. Chromatography  1.8.5. Centrifugation  1.8.6. Sublimation  1.8.7. Decantation and sedimentation
  2. Atomic Structure  2.1. Discovery of atoms  2.2. Dalton's atomic theory  2.3. Structure of the atom  2.4. Subatomic particles  2.5. Atomic number and mass number  2.6. Isotopes and Isobars  2.7. Electronic configuration  2.8. Bohr's atomic model  2.9. Modern Atomic Model (Quantum Mechanical Model - Introduction)  2.10. Valency and chemical bonding
  3. C hemical reactions and equations  3.1. Introduction to Chemical Reactions  3.2. Chemical Equation Formulation  3.3. Balancing Chemical Equations  3.4. Types of Chemical Reactions  3.4.1. Combination Reactions  3.4.2. Decomposition reactions  3.4.3. Displacement reactions  3.4.4. Double displacement reactions  3.4.5. Redox reactions  3.4.6. Endothermic and Exothermic Reactions  3.5. Effects of chemical reactions  3.6. Energy Changes in Chemical Reactions  3.7. Catalysts and their role in reactions
  4. Periodic table and periodicity  4.1. History of Periodic Classification  4.2. Mendeleev's periodic table  4.3. Modern Periodic Table  4.4. Periods and groups in the periodic table and periodicity  4.5. Trends in the Periodic Table  4.5.1. Atomic size  4.5.2. Ionization Energy  4.5.3. Electron affinity  4.5.4. Electronegativity  4.5.5. Metallic and Non-Metallic Properties  4.6. Noble gases and their properties
  5. Chemical bond  5.1. Introduction to Chemical Bonding  5.2. Types of chemical bonds  5.2.1. Ionic bond  5.2.2. Covalent bond  5.2.3. Metallic bond  5.2.4. Hydrogen bonding  5.2.5. Van der Waals force  5.3. Properties of Ionic and Covalent Compounds  5.4. Molecular Structure and Geometry (VSEPR Theory - Basics)
  6. Acids, Bases and Salts  6.1. Introduction to Acids and Bases  6.2. Properties of Acids  6.3. Properties of bases  6.4. pH scale and its importance  6.5. Indicators and their uses  6.6. Neutralization reactions  6.7. Strength of Acids and Bases (Strong vs. Weak)  6.8. Preparation of salts  6.9. Uses of acids, bases and salts in daily life
  7. Metals and Nonmetals  7.1. Physical Properties of Metals  7.2. Physical Properties of Nonmetals  7.3. Chemical properties of metals  7.4. Chemical Properties of Nonmetals  7.5. Reactivity Series of Metals  7.6. Extraction of metals  7.7. Corrosion and its prevention  7.8. uses of metals and nonmetals
  8. Carbon and its compounds  8.1. Introduction to Carbon  8.2. Allotropes of carbon (diamond, graphite, fullerene)  8.3. Hydrocarbons  8.3.1. Hydrocarbons  8.3.2. Alkene  8.3.3. Alkynes  8.4. Functional groups in organic chemistry  8.5. Isomerism in organic compounds  8.6. Alcohols and Carboxylic Acids  8.7. Soaps and detergents  8.8. Polymers (Introduction to Natural and Synthetic Polymers)
  9. Solutions and Mixtures  9.1. Types of mixtures  9.2. Homogeneous and Heterogeneous Mixtures  9.3. Concentration of solutions  9.4. Solubility and factors affecting solubility  9.5. Suspensions and Colloids  9.6. Saturated, unsaturated and supersaturated solutions
  10. Air and atmosphere  10.1. Composition of air  10.2. Role of gases in the atmosphere  10.4. Acid rain and its effects  10.5. Greenhouse effect and global warming  10.6. Ozone layer and its depletion  10.7. Air pollution control measures
  11. Water and its importance  11.1. Composition and properties of water  11.2. Hardness of water (temporary and permanent hardness)  11.3. Causes of water pollution  11.4. Effects of Water Pollution  11.5. Water purification methods (filtration, boiling, chlorination)
  12. Industrial Chemistry  12.1. Introduction to Industrial Chemistry  12.2. Important industrial chemicals (sulfuric acid, ammonia, sodium hydroxide)  12.3. Chemical processes in industry  12.4. Uses of Industrial Chemistry in Daily Life  12.5. Environmental impact of industrial chemical processes

Grade 9 → C hemical reactions and equations → Types of Chemical Reactions ↓


Redox reactions


Introduction to redox reactions

Redox reactions are chemical reactions in which the oxidation states of atoms change. These types of reactions occur through two main processes: oxidation and reduction. In simple terms, oxidation involves the loss of electrons, while reduction involves the gain of electrons. Redox reactions are essential to understanding many chemical processes, including biological systems, industrial applications, and environmental science.

Understanding oxidation and reduction

To fully understand redox reactions, we need to understand the concepts of oxidation and reduction:

  • Oxidation: It refers to the loss of electrons by a molecule, atom, or ion. This process increases the oxidation state.
  • Reduction: It refers to the gain of electrons by a molecule, atom, or ion. This process lowers the oxidation state.

Example of oxidation and reduction

Consider the reaction between hydrogen gas and fluorine gas to form hydrogen fluoride:

    H 2 + F 2 ⟶ 2HF

In this reaction, hydrogen is oxidized upon losing electrons, while fluorine is reduced upon gaining electrons.

The concept of oxidation number

Oxidation numbers help understand and balance redox reactions. The oxidation number is the theoretical charge an atom would have if the compound was made up of ions. Here are some rules for determining the oxidation number:

  • The oxidation number of an atom in its elemental form is 0.
  • The oxidation number of a monatomic ion is equal to its charge.
  • In compounds, the oxidation number of hydrogen is normally +1, while the oxidation number of oxygen is normally -2.
  • In a neutral compound the sum of the oxidation numbers is 0, while in a polyatomic ion it is equal to the charge of the ion.

Determination of oxidation and reduction

In a redox reaction, the change in oxidation number can be analyzed to determine which species has been oxidized and which has been reduced:

Example of identifying redox changes

Consider the reaction between iron and copper (II) sulfate:

    2H2O + CuSO4 ⟶ FeSO4 + Cu

Calculate the change in oxidation number:

  • Iron: 0 (in Fe) to +2 (in FeSO 4 ) - oxidized
  • Copper: +2 (in CuSO4 ) to 0 (in Cu) - less

Thus, iron gets oxidised and copper gets reduced.

Balancing redox reactions

Balancing redox reactions requires making sure that both the mass and charge are balanced. There are two common methods for balancing redox reactions: the oxidation number method and the half-reaction method.

Half-reaction method

This method involves splitting the redox reaction into two half-reactions: one for oxidation and one for reduction. Each half-reaction is balanced for mass and charge, and then they are combined to give a balanced redox equation.

Example of half-reaction method

Balance the redox reaction between zinc and hydrochloric acid:

    Zn + HCl ⟶ ZnCl 2 + H 2

Divided into half-reactions:

Oxidation: Zn ⟶ Zn 2+ + 2e -

Reduction: 2H + + 2e - ⟶ H 2

Combine to balance the reaction:

    2H2O + 2HCl ⟶ ZnCl2 + H2

Oxidation number method

This method assigns oxidation numbers to determine which elements are oxidized and which are reduced, and then uses these numbers to determine the electron transfer.

Example of oxidation number method

Balance the reaction between potassium dichromate and iron (II) sulfate in acidic solution:

    K 2 Cr 2 O 7 + FeSO 4 + H 2 SO 4 ⟶ Cr 2 (SO 4 ) 3 + Fe 2 (SO 4 ) 3 + H 2 O + K 2 SO 4

Determine the change in oxidation state:

  • Chromium: +6 to +3 (low)
  • Iron: +2 to +3 (oxidized)

Balance and combine the electron transfers:

Final Balanced Equation:

    K 2 Cr 2 O 7 + 6FeSO 4 + 7H 2 SO 4 ⟶ Cr 2 (SO 4 ) 3 + 3Fe 2 (SO 4 ) 3 + 7H 2 O + K 2 SO 4

Types of redox reactions

There are different types of redox reactions, which are generally classified based on their applications and properties:

Combination reactions

In these reactions, two or more substances combine to form a single product. In redox terms, one of the reactants is oxidized while the other is reduced.

Example of a combination reaction

Combination of Hydrogen Gas and Oxygen Gas:

    2H2 + O2 ⟶ 2H2O

In this reaction hydrogen is oxidised and oxygen is reduced.

Decomposition reactions

In these reactions, a single compound breaks down into two or more simpler products, often involving redox processes.

Example of a decomposition reaction

Decomposition of potassium chlorate:

    2KClO 3 ⟶ 2KCl + 3O 2

Here the chlorine is reduced, and oxygen is liberated.

Displacement reactions

These involve replacing one element in a compound with another. They often involve metals and are common redox reactions.

Example of a displacement reaction

Reaction between zinc and copper (II) sulfate:

    Zn + CuSO4 ⟶ ZnSO4 + Cu

Zinc displaces copper from the sulfate, thereby oxidizing the zinc and reducing the copper.

Tips for understanding redox reactions

  • Focus on changing oxidation states to identify redox pairs.
  • Practice balancing equations, especially using the half-reaction method for clarity.
  • Remember that oxidizing agents are reduced, and reducing agents are oxidized.
  • Focus on common redox reactions and their role in various scientific and industrial processes.

Applications of redox reactions

Redox reactions have wide applications in various fields:

In biological systems

Redox reactions are central to biological processes such as cellular respiration and photosynthesis. These reactions allow for the transfer and storage of energy necessary for life.

Cellular respiration

Glucose is oxidized to produce energy for the cells:

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

In the industry

Industries use redox reactions in the production of chemicals such as metals, fertilizers, and sulfuric acid.

Extraction of metals

Iron is extracted from its ore via the redox reaction:

    Fe 2 O 3 + 3CO ⟶ 2Fe + 3CO 2

In environmental science

Redox reactions are important in processes such as the breakdown of pollutants and the cycling of natural elements such as nitrogen and carbon.

Treatment of waste

Pollutants are broken down through redox reactions, making wastewater treatment efficient.

Conclusion

Redox reactions are an essential concept in chemistry as they are involved in both theoretical and practical applications. Understanding these reactions provides deep insights into various processes that are a part of everyday life as well as advanced scientific fields. Practicing and demonstrating various scenarios of redox reactions will build a strong foundation in chemistry that will be useful for both students and professionals in various disciplines alike.


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Redox reactions

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