Redox & Disproportionation

Assigning Oxidation States

• Elemental Rule: Atoms in their pure elemental form (e.g., $O_2$, $P_4$, $Na$) always have an oxidation state of $0$.
• Monatomic Ions: The oxidation state is equal to the charge of the ion (e.g., $Cl^-$ is $-1$, $Mg^{2+}$ is $+2$).
• Oxygen and Hydrogen: Oxygen is typically $-2$ (except in peroxides where it is $-1$), and Hydrogen is typically $+1$ (except in metal hydrides where it is $-1$).

Balancing Redox Equations

• Half-Reaction Method: Split the reaction into an oxidation half and a reduction half, balance all atoms except $O$ and $H$, then balance $O$ using $H_2O$ and $H$ using $H^+$.
• Charge Balancing: Add electrons ($e^-$) to the side with the higher positive charge to ensure both sides of the half-reaction have the same net charge.
• Final Integration: Multiply the half-reactions by integers so that the number of electrons lost equals the number gained, then add them together and cancel common species.

• Disproportionation Mechanism: This occurs when a single element in an intermediate oxidation state reacts to form two products: one where the element is in a higher oxidation state and one where it is in a lower state.

• Thermodynamic Stability: Disproportionation typically happens when the intermediate state is less stable than the combination of the higher and lower states under specific conditions (e.g., pH or temperature).

• Comproportionation (Synproportionation): The reverse process where two different oxidation states of the same element react to form a single, intermediate oxidation state product.

• Identification: To identify disproportionation, track the oxidation state of every atom of a specific element; if one atom goes 'up' and another of the same element goes 'down' from the same starting material, it is disproportionation.

• Agent vs. Process: A common point of confusion is that the Oxidizing Agent is the substance being reduced, while the Reducing Agent is the substance being oxidized.

• The Zero-Sum Check: Always ensure that the sum of oxidation states in a neutral compound is $0$, and in a polyatomic ion, it equals the ion's charge; this is the fastest way to verify assigned numbers.

• Identify the 'Spectators': In complex equations, identify elements whose oxidation states do not change (often $O$ and $H$); focusing only on the changing elements simplifies the balancing process.

• Medium Matters: If a reaction occurs in a basic medium, balance it as if it were in an acidic medium first, then add $OH^-$ ions to both sides to neutralize $H^+$ and form $H_2O$.

• Sanity Check: After balancing, verify that the total charge on the reactant side exactly matches the total charge on the product side; if they differ, the electron count is likely incorrect.