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 ↓


Oxidation number and its applications


Introduction

In chemistry, redox reactions play an important role, involving the transfer of electrons between substances. To better understand these reactions, chemists use the concept of oxidation numbers. Oxidation numbers indicate the degree of oxidation of an atom in a compound. Understanding oxidation numbers can help us balance chemical equations, predict the outcome of reactions, and understand the transfer of electrons between different species.

What is the oxidation number?

The oxidation number of an atom is the charge that the atom would have if the compound was made up entirely of ions. It is a theoretical number that helps understand redox reactions. This number can be positive, negative, or zero.

Basic rules for assigning oxidation numbers

There are several rules for assigning oxidation numbers:

  1. Elemental state: The oxidation number of an atom in its elemental form is always zero. For example: 
    H2, O2, N2, P4, S8 (Oxidation Number = 0)
  2. Monoatomic ions: For a monatomic ion, the oxidation number is equal to the charge of the ion. For example: 
    Na+ (Oxidation Number = +1), Cl- (Oxidation Number = -1)
  3. Oxygen: The oxidation number of oxygen is usually -2 in compounds, but there are some exceptions. For example, in peroxides such as H2O2, the oxidation number of oxygen is -1.
  4. Hydrogen: The oxidation number of hydrogen is typically +1 when bonded to non-metals and -1 when bonded to metals. For example: 
    H2O (Oxidation Number of H = +1), NaH (Oxidation Number of H = -1)
  5. Alkali Metals (Group 1): The oxidation number in the compounds of these elements is always +1.
  6. Alkaline Earth Metals (Group 2): The oxidation number in the compounds of these elements is always +2.
  7. Halogens: These usually have an oxidation number of -1 in compounds, unless they are combined with an element with higher electronegativities.
  8. Sum of oxidation numbers: For a neutral compound, the sum of the oxidation numbers of all atoms is zero. For polyatomic ions, the sum is equal to the charge of the ion.

Visual example of assigning oxidation numbers

Let's see how oxidation numbers are assigned using a simple example: KMnO4.

K: +1 (alkali metal) Mn: +7 (calculated) O: -2 (typical oxidation state for oxygen) Compound: KMnO4 Equation: +1 + Mn + 4(-2) = 0 Solve: +1 + Mn - 8 = 0 Mn = +7 (charge on manganese)

Understanding redox reactions with oxidation numbers

Redox reactions are characterized by a change in the oxidation number of the species involved. Oxidation refers to an increase in oxidation number, while reduction refers to a decrease in oxidation number. These reactions always occur together in nature. When we look at a chemical reaction, identifying these changes can help us understand the flow of electrons and predict the reaction products.

Example of a redox reaction

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

H2 + F2 → 2HF

Determination of oxidation number:

  • Before the response:
    • H in H2: 0
    • in F F2: 0
  • After the response:
    • H in HF: +1
    • F in HF: -1

In this response:

  • The oxidation number of hydrogen changes from 0 to +1 (oxidation).
  • The oxidation number of fluorine changes from 0 to -1 (reduction).

Role of oxidation numbers in balancing redox reactions

The oxidation number is important for balancing redox reactions. Let's explore the process using the example of an acidic solution: the reaction between potassium dichromate and iron (II) sulfate.

Unbalanced Reaction:

Cr2O72- + Fe2+ + H+ → Cr3+ + Fe3+ + H2O

Steps to balance:

  1. Recognize oxidation and reduction.
    • Cr2O72- (Cr changes from +6 to +3)
    • Conversion of Fe 2+ to Fe3+
  2. Equalize electron transfer.

    To balance this equation it is ensured that the number of electrons lost in oxidation is equal to the number of electrons gained in reduction.

Applications of oxidation numbers in real life

Beyond the theoretical implications, oxidation numbers are important in a variety of practical applications:

Predicting the feasibility of a response

By calculating oxidation numbers, we can sometimes predict whether a reaction is possible. If a substance does not undergo oxidation or reduction, a redox reaction is unlikely.

Understanding corrosion processes

Oxidation numbers help us understand corrosion, such as the rusting of iron. The oxidation number of iron increases as it turns into iron oxide (rust), which teaches us about prevention techniques using coatings or sacrificial anodes.

Analysis of biological processes

In biological systems, redox reactions are important in processes such as cellular respiration and photosynthesis, where substrates undergo complex electron transfers with different oxidation states.

Conclusion

Oxidation numbers serve as a theoretical bridging tool for understanding electron transfer processes in redox reactions. By knowing oxidation number assignments and rules, we can analyze complex chemical reactions and predict how atoms will interact in different scenarios. Whether balancing chemical equations or explaining biochemical phenomena, these numbers reveal the unseen world of chemical changes.


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Oxidation number and its applications

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