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


Redox reactions, also called reduction-oxidation reactions, are a family of reactions that involve the transfer of electrons between chemicals. These reactions are central to many biological and industrial processes, so understanding them is an essential part of chemistry.

What are redox reactions?

Redox reactions involve two major processes: oxidation and reduction.

  • Oxidation is the process in which a substance loses electrons.
  • Reduction is the process in which a substance gains electrons.

A mnemonic to help remember this is OIL RIG which stands for:

  • Oxidation Is Loose
  • Deduction is a benefit

Basic example of a redox reaction

A simple example of a redox reaction is the reaction between hydrogen and oxygen to form water:

2 H2 + O2 → 2 H2O

In this response:

  • Hydrogen (H2) is oxidized as it loses electrons. The oxidation state of hydrogen increases from 0 to +1.
  • Oxygen (O2) is reduced as it gains electrons. The oxidation state of oxygen decreases from 0 to -2.
H2 O2

The diagram above shows hydrogen and oxygen forming water through a redox reaction. Hydrogen atoms transfer electrons to oxygen atoms, forming water molecules.

Oxidation number

To fully understand redox reactions, we need to talk about oxidation numbers. Oxidation numbers are a way of keeping track of electrons during a chemical reaction. They can be thought of as imaginary charges assigned to atoms in molecules or ions.

Here are some basic rules for determining oxidation numbers:

  • The oxidation number of an atom in its elemental form is always zero. For example, the oxidation number of O2, H2, N2, etc. is zero.
  • For a simple (monatomic) ion, the oxidation number is equal to the charge on the ion. For example, the oxidation number of Na+ is +1, and the oxidation number of Cl- is -1.
  • The oxidation number of oxygen is usually -2, except when bonded to peroxides or fluorine.
  • The oxidation number of hydrogen is generally +1 when it combines with non-metals, and -1 when it combines with metals.
  • In a neutral compound the sum of the oxidation numbers must be zero.
  • The sum of the oxidation numbers in a polyatomic ion must be equal to the charge of the ion.

Let's apply these rules to find the oxidation number of water (H2O):

  • The oxidation number of oxygen is usually -2.
  • The oxidation number of each hydrogen is +1.
  • The sum of oxidation numbers in H2O = 2(+1) + (-2) = 0, which satisfies the rule of neutral compounds.

Identifying redox reactions

Identifying redox reactions involves checking for changes in the oxidation numbers of the substances in the reaction. If the oxidation numbers change, the reaction is redox.

Consider the following reaction:

4Zn + CuSO4 → ZnSO4 + Cu

To determine if this is a redox reaction, calculate the oxidation number:

  • In Zn: Oxidation number of Zn = 0, ZnSO4: Zn = +2 (in compound form).
  • In CuSO4: Cu = +2, elemental Cu: Cu = 0.

Information:

  • Zn goes from 0 to +2 (loses electrons, i.e. is oxidized).
  • Cu changes from +2 to 0 (gains electrons, i.e., is reduced).
Zinc CuSO4 ZnSO4 Cu

The diagram shows the transfer of electrons from zinc to copper, forming zinc sulfate and copper metal. Zinc oxidation forms zinc ions and in parallel copper gains electrons, turning it into native copper.

Balancing redox reactions

Balancing redox reactions ensures that the number of electrons lost in oxidation is equal to the number of electrons gained in reduction. Here's how to balance using the half-reaction method:

Steps to balancing redox reactions

  1. Split the unbalanced redox equation into two half-reactions: one for oxidation and one for reduction.
  2. Balance each half-reaction for mass, then charge using electrons.
  3. Multiply the half-reactions by appropriate integers so that the electrons in both half-reactions are equal.
  4. Recombine the balanced half-reactions, making sure that the electrons cancel out.
  5. Verify the atomic and charge balance in the last equation.

Let's balance the redox reaction of iron with chlorine:

2Fe + Cl2 → FeCl3

Balancing process

Step 1: Split into half-reactions:

Oxidation: Fe → Fe3+ + 3e-
Reduction: Cl2 + 2e- → 2Cl-

Step 2: Balance atoms other than O and H, then balance oxygen and hydrogen if present (not needed here).

Step 3: Balance the charges by adding electrons:

  • The oxidation is already balanced for the charge.
  • 3 electrons required for reduction: Cl2 + 6e- → 2Cl2

Step 4: Multiply the 2 reductions by appropriate integers to make the electrons equal to the oxidations:

Oxidation: Fe → Fe3+ + 3e-
Reduction: Cl2 + 6e- → 2Cl- (same electron gain required)

Step 5: Add and remove electrons:

2Fe + 3/2Cl2 → FeCl3

The balanced reaction ensures equality in electron transfer and atom list.

Common redox reactions in everyday life

Redox reactions are all around us and are part of everyday life. Here are some examples:

Combustion

Combustion is a redox reaction in which fuel reacts with oxygen to release energy. For example, the burning of methane gas:

CH4 + 2 O2 → CO2 + 2 H2O
CH4 2 O2 CO2 + 2 H2O

Batteries

Batteries operate on redox reactions, driven by chemical differences to transfer electrons from the anode to the cathode. In a simple zinc-carbon battery:

Zn → Zn2+ + 2e-
2MnO2 + 2e- + 2NH4Cl → Mn2O3 + H2O + 2NH3

War

Rusting of iron is a redox process in which iron is oxidized by oxygen and water to form iron oxide.

4 Fe + 3 O2 + 6 H2O → 4 Fe(OH)3

Conclusion

Redox reactions are fundamental to understanding both basic and complex chemical processes. From the burning of fuels to biological processes and energy storage in batteries, redox reactions facilitate essential transformations involving electron transfer. Understanding the basics of oxidation and reduction, as well as efficiently determining oxidation numbers, is critical to exploring the wide applications and implications of redox chemistry.


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Redox reactions

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