Redox Chemistry: Oxidation, Reduction, and Oxidation Numbers

Vocabulary-style flashcards covering the principles of redox reactions, oxidation number rules, and the methodology for balancing half-equations based on lecture notes.

Reducing agents

Electron donors that cause another species to be reduced while being oxidised themselves.

Oxidising agents

Electron acceptors that cause another species to be oxidised while being reduced themselves.

Oxidation

The process of electron loss (e.g., $Zn \rightarrow Zn^{2+} + 2e^{-}$) involving an increase in oxidation number.

Reduction

The process of electron gain (e.g., $Cl_2 + 2e^{-} \rightarrow 2Cl^{-}$) involving a decrease in oxidation number.

Oxidation Number of Uncombined Elements

The oxidation number is zero for all uncombined elements, such as $Zn$, $Cl_2$, $O_2$, and $Ar$.

Oxidation Number Sum (Compounds)

The oxidation numbers of all elements in a neutral compound must add up to zero.

Oxidation Number (Monoatomic Ion)

The oxidation number is equal to the ionic charge, such as $+2$ for $Zn^{2+}$ and $-1$ for $Cl^{-}$.

Oxidation Number Sum (Polyatomic Ions)

The sum of individual oxidation numbers of the elements adds up to the charge on the ion, such as $-2$ for $CO_3^{2-}$.

Invariable Oxidation Number (Group 1 Metals)

$+1$ in common compounds.

Invariable Oxidation Number (Group 2 Metals)

$+2$ in common compounds.

Invariable Oxidation Number (Aluminum)

$+3$ in common compounds.

Oxidation Number of Hydrogen

$+1$, except in metal hydrides (e.g., $NaH$) where it is $-1$.

Oxidation Number of Fluorine

$-1$ in all of its compounds.

Oxidation Number of Chlorine, Bromine, and Iodine

$-1$, except when they are in compounds with oxygen or fluorine.

Oxidation Number of Oxygen

$-2$, except in peroxides (e.g., $H_2O_2$) where it is $-1$, and in compounds with fluorine.

Reduction Half-Equation

A chemical equation showing only the reduction process, with electrons written on the left (reactant) side.

Oxidation Half-Equation

A chemical equation showing only the oxidation process, with electrons written on the right (product) side.

Balancing O in Complex Half-Equations

In acidic conditions, balance oxygen by adding $H_2O$ to the side lacking oxygen atoms.

Balancing H in Complex Half-Equations

In acidic conditions, balance hydrogen by adding $H^{+}$ ions to the side lacking hydrogen atoms.

Combining Half-Equations

The process of multiplying half-equations to ensure equal numbers of electrons so they cancel out when added together to form a full redox equation.