Redox Reactions — Revision Notes

⚡ 30-Second Revision

Oxidation: — Loss of $e^-$, Oxidation State (OS) increases.
Reduction: — Gain of $e^-$, OS decreases.
Oxidizing Agent: — Gets reduced, causes oxidation.
Reducing Agent: — Gets oxidized, causes reduction.
OS Rules:

- Element: 0 - Monatomic ion: Charge - Group 1: +1; Group 2: +2 - H: +1 (non-metals), -1 (metal hydrides) - O: -2 (most), -1 (peroxides), -1/2 (superoxides), +2 ($OF_2$) - Sum of OS = 0 (neutral compound) or ion charge (polyatomic ion).

Balancing Methods: — Ion-electron (half-reaction) method, Oxidation number method.
Acidic Medium: — Balance O with $H_2O$, H with $H^+$.
Basic Medium: — Balance O with $H_2O$, H with $H_2O$ and $OH^-$.
Disproportionation: — Same element oxidized and reduced.

2-Minute Revision

Redox reactions are defined by the transfer of electrons. Oxidation is the loss of electrons, leading to an increase in oxidation state, while reduction is the gain of electrons, causing a decrease in oxidation state.

These processes are always coupled. The species that gets oxidized acts as the reducing agent, and the species that gets reduced acts as the oxidizing agent. Mastering oxidation state assignment is crucial; remember the specific rules for common elements like oxygen (e.

g., -1 in peroxides) and hydrogen (e.g., -1 in metal hydrides), and that the sum of oxidation states equals the compound's charge. Balancing redox reactions is a key skill, primarily using the ion-electron method.

For acidic media, balance oxygen with $H_2O$ and hydrogen with $H^+$. For basic media, balance oxygen with $H_2O$ and hydrogen with $H_2O$ on the H-deficient side and $OH^-$ on the opposite side. Pay special attention to disproportionation reactions where a single element is both oxidized and reduced.

5-Minute Revision

Redox reactions are the backbone of many chemical transformations, characterized by electron transfer. Oxidation is the process where a species loses electrons, resulting in an increase in its oxidation state (e.

g., $Fe^{2+} \rightarrow Fe^{3+} + e^-$). Conversely, reduction is the gain of electrons, leading to a decrease in oxidation state (e.g., $Cu^{2+} + 2e^- \rightarrow Cu$). These two processes are inseparable.

The species that undergoes oxidation is the reducing agent, and the species that undergoes reduction is the oxidizing agent.

Assigning oxidation states accurately is fundamental. Recall the rules: elements are 0; monatomic ions equal their charge; Group 1 metals are +1, Group 2 are +2; hydrogen is +1 (except in metal hydrides, -1); oxygen is -2 (except in peroxides -1, superoxides -1/2, $OF_2$ +2). The sum of oxidation states in a neutral compound is zero, and in an ion, it equals the ion's charge. For example, in $KMnO_4$, Mn is +7.

Balancing redox reactions is a critical skill for NEET. The ion-electron (half-reaction) method is preferred. Separate the reaction into oxidation and reduction half-reactions. Balance atoms (non-O/H first, then O with $H_2O$, then H with $H^+$ in acidic medium or $H_2O$/$OH^-$ in basic medium).

Finally, balance charge by adding electrons. Multiply half-reactions to equalize electron transfer, then combine. For instance, balancing $MnO_4^- + Fe^{2+} \rightarrow Mn^{2+} + Fe^{3+}$ in acidic medium requires $5Fe^{2+}$ and $2MnO_4^-$.

Remember special types like disproportionation reactions, where one element is simultaneously oxidized and reduced (e.g., $H_2O_2 \rightarrow H_2O + O_2$). A solid grasp of these concepts and consistent practice with balancing problems will ensure success in this high-scoring topic.

Prelims Revision Notes

Redox Reactions: NEET Quick Recall

1. Definitions:

Oxidation: — Loss of electrons (LEO), increase in oxidation state (OS).
Reduction: — Gain of electrons (GER), decrease in oxidation state (OS).
Oxidizing Agent (Oxidant): — Causes oxidation, itself gets reduced. OS decreases.
Reducing Agent (Reductant): — Causes reduction, itself gets oxidized. OS increases.

2. Rules for Assigning Oxidation States (OS):

Elemental form: — OS = 0 (e.g., $O_2$, $Na$, $P_4$).
Monatomic ions: — OS = charge (e.g., $Na^+$ is +1, $Cl^-$ is -1).
Group 1 metals: — Always +1 in compounds.
Group 2 metals: — Always +2 in compounds.
Hydrogen (H): — +1 with non-metals (e.g., $H_2O$), -1 with metals (e.g., $NaH$).
Oxygen (O): — Usually -2. Exceptions:

* Peroxides ($H_2O_2$, $Na_2O_2$): -1 * Superoxides ($KO_2$): -1/2 * With Fluorine ($OF_2$): +2

Fluorine (F): — Always -1.
Sum of OS: — 0 for neutral compounds; equals ion charge for polyatomic ions.

* *Example:* OS of Cr in $Cr_2O_7^{2-}$: $2x + 7(-2) = -2 Rightarrow 2x - 14 = -2 Rightarrow 2x = 12 Rightarrow x = +6$.

3. Balancing Redox Reactions (Ion-Electron Method):

Step 1: — Separate into two half-reactions (oxidation and reduction).
Step 2: — Balance atoms other than O and H.
Step 3 (Acidic Medium):

* Balance O by adding $H_2O$ to the side deficient in O. * Balance H by adding $H^+$ to the side deficient in H.

Step 3 (Basic Medium):

* Balance O by adding $H_2O$ to the side deficient in O. * Balance H by adding $H_2O$ to the side deficient in H, and an equal number of $OH^-$ to the opposite side.

Step 4: — Balance charge by adding electrons ($e^-$) to the more positive side.
Step 5: — Equalize electrons in both half-reactions by multiplying by appropriate integers.
Step 6: — Add the two half-reactions and cancel common species ($e^-$, $H_2O$, $H^+$/$OH^-$).

4. Types of Redox Reactions:

Combination: — $A + B \rightarrow C$
Decomposition: — $A \rightarrow B + C$
Displacement: — $A + BC \rightarrow AC + B$
Disproportionation: — Same element is both oxidized and reduced (e.g., $2H_2O_2 \rightarrow 2H_2O + O_2$).

5. Key Oxidizing/Reducing Agents:

Oxidizing: — $KMnO_4$ ($Mn^{+7} \rightarrow Mn^{+2}$ in acid), $K_2Cr_2O_7$ ($Cr^{+6} \rightarrow Cr^{+3}$ in acid), $HNO_3$, $H_2SO_4$ (conc.), halogens.
Reducing: — Alkali metals, $H_2$, $C$, $CO$, $SnCl_2$, $H_2S$, $SO_2$, $HI$, $HBr$.

Vyyuha Quick Recall

OIL RIG

Oxidation Is Loss (of electrons) Reduction Is Gain (of electrons)