Electrolysis, particularly of aqueous solutions, requires a significant amount of energy mainly in the form of electricity. When metals are obtained from their ores using electrolysis, the energy costs can be high due to the required voltage and the duration of the process. Reduction methods, on the other hand, such as the reduction of metal oxides with carbon in a blast furnace, often utilise chemical reactions that release energy. These exothermic reactions can aid in the overall process, making it more energy-efficient. Moreover, the raw materials like coke used in blast furnaces for reduction are often cheaper than the electricity costs associated with electrolysis.

Overpotential refers to the extra voltage (over and above the theoretical voltage) required during electrolysis to drive the reaction at a noticeable rate. The nature of the electrode material can influence this overpotential. Inert electrodes like platinum or graphite have higher overpotentials for the evolution of oxygen or hydrogen as compared to more reactive metals. The material of the electrode can influence the ease with which a reaction takes place on its surface. A metal electrode that can form an oxide layer, for instance, may show a higher overpotential for oxygen evolution because the oxide layer hinders the process.

In a molten state, the ionic compound is purely composed of its constituent ions without the presence of water. Therefore, water doesn't compete in the redox processes during electrolysis. For example, in the molten state electrolysis of sodium chloride, sodium ions are reduced at the cathode to produce sodium metal, while chloride ions are oxidised at the anode to produce chlorine gas. However, in an aqueous solution of sodium chloride, as discussed earlier, water interferes, leading to the evolution of hydrogen gas at the cathode instead of sodium metal deposition.

Pure water is a poor conductor of electricity due to its very low concentration of ions. For electrolysis to occur spontaneously, there must be sufficient ions present to carry the current. Additionally, the reduction potential for hydrogen ions (from water) to form hydrogen gas and the oxidation potential for water to form oxygen gas are not favourable enough to occur spontaneously. To drive the non-spontaneous electrolysis of pure water, an external voltage (greater than the thermodynamic threshold) is required. In practical scenarios, to aid the electrolysis process, an electrolyte is added to increase the conductivity and reduce the voltage required.

Sodium is a highly reactive metal, and in an aqueous environment like that of a sodium chloride solution, water is also present as a potential reactant. During the electrolysis of an aqueous sodium chloride solution, there are two potential cations that can be reduced at the cathode: sodium ions (Na+) and hydrogen ions (H+ from water). However, sodium is more reactive than hydrogen, meaning it has a greater tendency to lose electrons (get oxidised) than to gain them (get reduced). As a result, hydrogen ions are preferentially reduced at the cathode, producing hydrogen gas, while the sodium ions remain in the solution.