Ncert Books Class 11 Chemistry Chapter 8 Free Download

To determine the oxidation number, you apply a set of hierarchical rules. For example, the oxidation state of an element in its free form is zero, and for a simple ion, it is its charge. However, some important exceptions frequently tested are:

• Oxygen: Usually -2, but it is -1 in peroxides (e.g., H₂O₂), -1/2 in superoxides (e.g., KO₂), and +2 in oxygen difluoride (OF₂).
• Hydrogen: Usually +1, but it is -1 in metal hydrides (e.g., NaH, CaH₂).
• Halogens: Usually -1, but they can have positive oxidation states when bonded to a more electronegative atom, like oxygen (e.g., in chlorates and perchlorates).

This is a fundamental concept. A redox reaction is treated as a single process because oxidation and reduction are interdependent and must occur at the same time. Oxidation involves the loss of electrons, while reduction involves the gain of electrons. An electron cannot be simply 'lost' into space; if one substance loses electrons (gets oxidised), another substance must be present to accept those electrons (get reduced). Therefore, the two half-reactions always happen together, making the overall redox reaction a single, unified electron transfer event.

A disproportionation reaction is a specific type of redox reaction where an element in a particular oxidation state is simultaneously oxidised and reduced. This is an important concept for HOTS (Higher Order Thinking Skills) questions. A classic example asked in exams is the reaction of phosphorus with aqueous alkali:
P₄(s) + 3OH⁻(aq) + 3H₂O(l) → PH₃(g) + 3H₂PO₂⁻(aq)
Here, the oxidation state of phosphorus (0) changes to -3 in phosphine (PH₃) (reduction) and to +1 in the hypophosphite ion (H₂PO₂⁻) (oxidation).

The most reliable method is by tracking the change in oxidation numbers. First, assign oxidation numbers to all elements in the reactants and products.

• The substance containing the element whose oxidation number increases is the reducing agent (because it itself gets oxidised).
• The substance containing the element whose oxidation number decreases is the oxidising agent (because it itself gets reduced).

Both methods are valid for balancing redox reactions, but they differ in their approach.

• The Oxidation Number Method focuses on equalising the total increase and decrease in oxidation numbers for the atoms involved.
• The Half-Reaction Method splits the reaction into two separate half-equations (one for oxidation, one for reduction), balances the atoms and charges in each, and then combines them.

For the Class 11 final exams, you should be able to identify and provide examples for the following four types of redox reactions:

• Combination Reactions: Where two or more substances combine to form a single compound, e.g., C(s) + O₂(g) → CO₂(g).
• Decomposition Reactions: Where a compound breaks down into two or more components, e.g., 2H₂O(l) → 2H₂(g) + O₂(g).
• Displacement Reactions: Where an ion or atom in a compound is replaced by an ion or atom of another element, e.g., Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s).
• Disproportionation Reactions: Where the same element is both oxidised and reduced, e.g., 2H₂O₂(aq) → 2H₂O(l) + O₂(g).

A solid foundation in redox reactions is critical for advanced chemistry topics. Its importance extends to:

• Electrochemistry: Understanding how batteries, galvanic cells, and electrolytic cells work is entirely based on redox principles.
• Metallurgy: The extraction of metals from their ores (e.g., iron from haematite) involves key reduction processes.
• Biochemistry: Cellular respiration, the process that provides energy to living organisms, is a complex series of redox reactions.
• Analytical Chemistry: Many quantitative analysis techniques, like redox titrations, are used to determine the concentration of substances.