1. Assign oxidation numbers. 2. Write half-reactions. 3. Balance elements (except O and H). 4. Balance O with H₂O. 5. Balance H with H⁺. 6. Balance charge with e⁻. 7. Combine half-reactions.

1. Balance as in acidic solution. 2. Neutralize H⁺ with OH⁻ to form H₂O. 3. Add the same amount of OH⁻ to the other side. 4. Simplify by canceling H₂O molecules.

1. Assign oxidation numbers. 2. Write half-reactions. 3. Balance elements (except O and H). 4. Balance oxygen atoms by adding H₂O molecules.

5. Balance hydrogen atoms by adding H⁺ ions. 6. Balance the charge by adding electrons. 7. Combine the half-reactions, canceling out any common species.

8. For basic solutions: Neutralize H⁺ ions with OH⁻ to form H₂O, and add the same number of OH⁻ to the other side. 9. Double-check that atoms and charges are balanced.

The atom undergoes oxidation, and its oxidation number increases.

The atom undergoes reduction, and its oxidation number decreases.

Water ($H_2O$) is formed, and you must add the same amount of $OH^−$ to the other side of the equation to maintain balance.

The overall equation will not be balanced, leading to incorrect stoichiometric coefficients and an inaccurate representation of the reaction.

$Cr_2O_7^{2−}$ is reduced to $Cr^{3+}$, and $I^−$ is oxidized to $I_2$.

A reaction involving the transfer of electrons between atoms, leading to changes in oxidation states.

An atom's charge in a compound, showing how many electrons it has gained or lost compared to its neutral state.

The loss of electrons by an atom, resulting in an increase in its oxidation number.

The gain of electrons by an atom, resulting in a decrease in its oxidation number.

A representation showing either the oxidation or reduction process separately, including electrons.