Electrolysis of Aqueous Solutions

• Standard Reduction Potentials ($E^0$): The species with the most positive reduction potential is the most likely to be reduced at the cathode. Conversely, the species with the most negative reduction potential (or the most positive oxidation potential) is the most likely to be oxidized at the anode.

• Water Reduction: At the cathode, water can be reduced to hydrogen gas and hydroxide ions: $2H_2O(l) + 2e^- \rightarrow H_2(g) + 2OH^-(aq)$. This reaction has a standard potential of $-0.83V$, which competes with the reduction of metal cations.

• Water Oxidation: At the anode, water can be oxidized to oxygen gas and hydrogen ions: $2H_2O(l) \rightarrow O_2(g) + 4H^+(aq) + 4e^-$. This reaction has a standard potential of $+1.23V$, competing with the oxidation of anions.

Predicting Cathode Products

• Rule of Reactivity: Compare the metal cation to Hydrogen in the reactivity series. If the metal is more reactive than hydrogen (e.g., $Na^+, K^+, Mg^{2+}$), water is reduced and Hydrogen gas is evolved. If the metal is less reactive (e.g., $Cu^{2+}, Ag^+$), the metal itself is deposited on the electrode.

Predicting Anode Products

• Halide Concentration: If the solution contains a high concentration of halide ions ($Cl^-, Br^-, I^-$), the halogen gas is produced because halides are easier to oxidize than water under these conditions. If the anions are polyatomic (e.g., $SO_4^{2-}, NO_3^-$), they are too stable to be oxidized, and Oxygen gas is produced from water oxidation.

Step-by-Step Analysis

1. List all species present ($H_2O$, cations, anions).
2. Identify possible half-reactions at each electrode.
3. Apply reactivity rules for the cathode and concentration/stability rules for the anode.
4. Write the final half-equations and the overall cell reaction.

• Molten vs. Aqueous: In molten electrolysis, there is no competition from water, so the metal cation always forms the metal and the anion always forms the non-metal. In aqueous electrolysis, the presence of water often results in the production of $H_2$ or $O_2$ gas instead.

• The 'Big Four' Rule: Always remember that $K^+, Na^+, Li^+$, and $Ca^{2+}$ will NEVER be reduced in aqueous solution; they always result in $H_2$ gas. This is a very common exam shortcut for identifying cathode products.

• pH Changes: Be prepared to explain why the solution becomes acidic or basic. Reduction of water produces $OH^-$, increasing $pH$ at the cathode, while oxidation of water produces $H^+$, decreasing $pH$ at the anode.

• Inert vs. Active Electrodes: Always check if the electrodes are 'inert' (like graphite or platinum). If the anode is 'active' (like copper in a copper sulfate solution), the electrode itself may oxidize instead of the water or the anions.

• Predicting Sodium Metal: A frequent error is predicting that sodium metal forms during the electrolysis of aqueous $NaCl$. Because sodium is highly reactive, it would immediately react with water even if it did form; thus, hydrogen gas is the only logical product.

• Sulfate Oxidation: Students often try to oxidize sulfate ($SO_4^{2-}$) or nitrate ($NO_3^-$) ions. These ions are in their highest oxidation states and are extremely stable; in aqueous solutions, water will always be oxidized in their place.

• Overpotential: Theoretically, water oxidation requires $+1.23V$, but in practice, it often requires a higher voltage (overpotential) to produce gas. This explains why concentrated chloride ions are oxidized preferentially over water despite similar potentials.