Reactions at Electrodes

The anode is the positive electrode that attracts negatively charged ions, known as anions, from the electrolyte solution or melt.

Upon contact with the anode, anions surrender their extra electrons to the electrode; these electrons then flow through the external circuit toward the power source.

In aqueous solutions, if halide ions ($Cl^-$, $Br^-$, $I^-$) are present in high concentrations, they are discharged to form halogen gases (e.g., $2Cl^- \rightarrow Cl_2 + 2e^-$).

If no halides are present or the solution is dilute, hydroxide ions ($OH^-$) from water are discharged instead, producing oxygen gas and water: $4OH^- \rightarrow O_2 + 2H_2O + 4e^-$.

The cathode is the negative electrode that attracts positively charged ions, known as cations, which migrate through the electrolyte.

At the cathode surface, cations accept electrons provided by the external circuit to become neutral metal atoms or hydrogen gas molecules.

In aqueous electrolysis, the reactivity series determines the product: the least reactive cation is preferentially discharged.

If the metal in the salt is more reactive than hydrogen (e.g., Sodium, Magnesium), hydrogen gas is produced from $H^+$ ions: $2H^+ + 2e^- \rightarrow H_2$.

Half-equations are used to represent the specific chemical changes occurring at each individual electrode, focusing on the movement of electrons.

To construct a half-equation, first write the symbol for the ion and the resulting neutral product, then balance the number of atoms on both sides.

Finally, add electrons ($e^-$) to one side of the equation to ensure the total electrical charge is balanced (e.g., $Al^{3+} + 3e^- \rightarrow Al$).

It is critical to remember that electrons appear on the right side for oxidation (loss) and on the left side for reduction (gain).

In molten electrolysis, only the ions from the ionic compound itself are present, meaning the metal always forms at the cathode and the non-metal at the anode.

In aqueous electrolysis, water molecules partially dissociate into $H^+$ and $OH^-$ ions, creating competition for discharge at both electrodes.

Check the State: Always identify if the electrolyte is (l) for liquid/molten or (aq) for aqueous, as this completely changes the predicted products.

Balance Charges: Ensure the total charge on the left side of a half-equation equals the total charge on the right; a common mistake is placing electrons on the wrong side.

Identify Observations: Exams often ask for what you 'see'; remember that gases like $Cl_2$ (pale green) or $O_2$ (colorless bubbles) have distinct physical appearances.

Reactivity Check: For aqueous solutions, always compare the metal cation to hydrogen; if the metal is high in the reactivity series, expect $H_2$ gas at the cathode.