Concept of Oxidation and Reduction — Revision Notes

⚡ 30-Second Revision

Oxidation: — Loss of electrons; Increase in oxidation number.
Reduction: — Gain of electrons; Decrease in oxidation number.
Redox Reaction: — Oxidation and reduction occur simultaneously.
Oxidizing Agent: — Causes oxidation (gets reduced itself).
Reducing Agent: — Causes reduction (gets oxidized itself).
Oxidation Number Rules:

- Free element: 0 - Monatomic ion: equals its charge - Group 1: +1; Group 2: +2 - F: -1 - O: -2 (except in peroxides -1, superoxides -1/2, $OF_2$ +2) - H: +1 (except in metal hydrides -1) - Sum in neutral compound: 0 - Sum in polyatomic ion: equals ion's charge

Disproportionation: — Same element both oxidized and reduced.

2-Minute Revision

The core of redox chemistry lies in the transfer of electrons. Oxidation is defined as the loss of electrons, which always results in an *increase* in the oxidation number of the atom or ion. Conversely, reduction is the gain of electrons, leading to a *decrease* in its oxidation number.

These two processes are inseparable and always occur together in a redox reaction. The species that causes oxidation by accepting electrons is called the oxidizing agent (and itself gets reduced).

The species that causes reduction by donating electrons is the reducing agent (and itself gets oxidized). A crucial skill is to accurately assign oxidation numbers using a set of rules, paying attention to exceptions for oxygen and hydrogen.

A special type of redox reaction is disproportionation, where a single element in a reactant is simultaneously oxidized and reduced. Mastering these definitions and the ability to calculate oxidation numbers are essential for identifying redox reactions, their agents, and for subsequent topics like balancing equations and electrochemistry.

5-Minute Revision

To thoroughly revise oxidation and reduction, focus on the interconnected definitions and practical application. Start with the modern electronic definitions: **Oxidation is the loss of electrons, leading to an increase in oxidation number.

** For example, $Na \rightarrow Na^+ + e^-$, where Na's oxidation number goes from 0 to +1. Reduction is the gain of electrons, leading to a decrease in oxidation number. For instance, $Cl_2 + 2e^- \rightarrow 2Cl^-$, where Cl's oxidation number goes from 0 to -1.

Remember, these processes are always coupled in a redox reaction.

Next, understand the roles of agents. An oxidizing agent is the electron acceptor; it *causes* oxidation in another substance and *gets reduced* itself. Common examples include $KMnO_4$ and $O_2$. A reducing agent is the electron donor; it *causes* reduction in another substance and *gets oxidized* itself. Examples include Na and $H_2$. A common mistake is confusing the process with the agent's role.

Master the rules for assigning oxidation numbers. Key rules include: free elements are 0; monatomic ions equal their charge; Group 1 metals are +1, Group 2 are +2; Fluorine is always -1. Oxygen is typically -2, but -1 in peroxides ($H_2O_2$) and +2 in $OF_2$.

Hydrogen is +1, but -1 in metal hydrides ($NaH$). The sum of oxidation numbers in a neutral compound is zero, and in a polyatomic ion, it equals the ion's charge. Practice calculating oxidation numbers for various compounds like $Cr_2O_7^{2-}$ (Cr is +6) or $S_2O_3^{2-}$ (S is +2).

Finally, be aware of disproportionation reactions, where one element in a single reactant is simultaneously oxidized and reduced. A classic example is $2H_2O_2 \rightarrow 2H_2O + O_2$, where oxygen (in -1 state in $H_2O_2$) goes to -2 in $H_2O$ and 0 in $O_2$. This comprehensive understanding will prepare you for diverse NEET questions.

Prelims Revision Notes

Oxidation and Reduction: Quick Recall for NEET

1. Core Definitions:

Oxidation (Electronic): — Loss of electrons. Always results in an increase in oxidation number (ON).

* Example: $Na \rightarrow Na^+ + e^-$ (ON: 0 to +1)

Reduction (Electronic): — Gain of electrons. Always results in a decrease in oxidation number (ON).

* Example: $Cl_2 + 2e^- \rightarrow 2Cl^-$ (ON: 0 to -1)

Redox Reaction: — Oxidation and reduction always occur simultaneously.

2. Agents:

Oxidizing Agent (Oxidant): — The substance that *causes* oxidation in another. It *accepts electrons* and therefore *gets reduced* itself. (ON decreases).

* Examples: $KMnO_4$, $K_2Cr_2O_7$, $O_2$, $HNO_3$.

Reducing Agent (Reductant): — The substance that *causes* reduction in another. It *donates electrons* and therefore *gets oxidized* itself. (ON increases).

* Examples: $Na$, $Mg$, $H_2$, $C$, $H_2S$, $SnCl_2$.

3. Rules for Assigning Oxidation Numbers (ON):

Free elements: — ON = 0 (e.g., $O_2$, $H_2$, $Fe$).
Monatomic ions: — ON = charge of the ion (e.g., $Na^+$ is +1, $Cl^-$ is -1).
Group 1 metals (Li, Na, K, etc.): — Always +1 in compounds.
Group 2 metals (Be, Mg, Ca, etc.): — Always +2 in compounds.
Fluorine (F): — Always -1 in compounds.
Oxygen (O): — Usually -2.

* Exceptions: Peroxides ($H_2O_2$, $Na_2O_2$) = -1; Superoxides ($KO_2$) = -1/2; In $OF_2$ = +2.

Hydrogen (H): — Usually +1.

* Exceptions: Metal hydrides ($NaH$, $CaH_2$) = -1.

Sum of ONs: — In a neutral compound, sum = 0. In a polyatomic ion, sum = charge of the ion.

4. Types of Redox Reactions:

Disproportionation Reaction: — A single element in a reactant is simultaneously oxidized and reduced.

* Example: $2H_2O_2 \rightarrow 2H_2O + O_2$ (Oxygen from -1 to -2 and 0).

5. Common Traps:

Confusing 'oxidized' with 'oxidizing agent'.
Incorrectly applying oxidation number rules, especially for exceptions.
Assuming all reactions involving oxygen/hydrogen are redox, or that all redox reactions must involve them.

Vyyuha Quick Recall

OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).