Electrolytic and galvanic cells are two fundamental types of electrochemical cells, each operating on different principles. In a galvanic cell, chemical energy is converted into electrical energy through spontaneous redox reactions. These cells consist of two different metals connected by a salt bridge or a porous partition, with each metal immersed in a solution containing its ions. The chemical reaction between the metals and their ions generates a flow of electrons, creating an electric current. The classic example is the Daniell cell, with zinc and copper electrodes in their respective sulfate solutions.

In contrast, an electrolytic cell uses electrical energy to drive a non-spontaneous chemical reaction. In this setup, an external power source, like a battery or power supply, is connected to two electrodes immersed in an electrolyte. The applied electric current causes ions in the electrolyte to undergo oxidation or reduction at the electrodes, leading to chemical changes. Electrolytic cells are used for various applications, including electroplating, refining metals, and electrolysis of water.

The key difference lies in their operation: galvanic cells generate electricity through spontaneous reactions, whereas electrolytic cells use electricity to induce chemical changes.

Hydrogen gas is often produced at the cathode in aqueous solutions during electrolysis due to the presence of water. Water molecules dissociate into hydrogen (H⁺) and hydroxide (OH^-) ions. In solutions of salts of more reactive metals (like sodium or calcium), these metal cations compete with hydrogen ions for electrons at the cathode. However, since these metals are more reactive than hydrogen, they have a greater tendency to remain as ions. As a result, the less reactive hydrogen ions are reduced in preference to the metal ions, leading to the production of hydrogen gas.

The production of hydrogen gas at the cathode is also influenced by factors like the electrode material, the concentration of the electrolyte, and the applied voltage. For instance, in a dilute solution, there is a higher concentration of water molecules, which increases the likelihood of hydrogen ion reduction. Furthermore, some electrode materials, like platinum, can catalyze the reduction of hydrogen ions, making the production of hydrogen gas more favorable. Understanding the conditions that favor the reduction of hydrogen ions is crucial in electrolysis, particularly in processes like water treatment and the production of hydrogen fuel.

Overvoltage, also known as overpotential, is a phenomenon in electrolysis where a higher voltage than the thermodynamically predicted value is required to drive an electrochemical reaction. This occurs due to various factors, including the nature of the electrode material, the concentration of ions, and the presence of impurities. Overvoltage affects electrolysis by making the process less efficient and more energy-intensive. For instance, in the electrolysis of water, hydrogen and oxygen gases are produced at the cathode and anode, respectively. However, due to overvoltage, the actual voltage required to initiate and sustain the electrolysis of water is higher than the theoretical value. This means more electrical energy is consumed than what would be expected based on standard electrode potentials alone. Overvoltage is particularly significant in industrial applications where energy efficiency directly impacts the cost and environmental impact of the process. Understanding and minimizing overvoltage is crucial for developing efficient and sustainable electrolytic systems.

The concentration of the electrolyte has a significant impact on the products of electrolysis, particularly in aqueous solutions. In dilute solutions, water molecules can compete with the dissolved ions for discharge at the electrodes. For instance, in a dilute solution of sodium chloride, there is a higher probability of water molecules being reduced at the cathode to produce hydrogen gas, rather than sodium ions being reduced. Conversely, in a concentrated solution of the same electrolyte, the higher concentration of sodium and chloride ions increases the likelihood of sodium being reduced at the cathode and chlorine being oxidized at the anode. Additionally, the concentration of the electrolyte influences the rate of electrolysis: a more concentrated solution typically leads to a higher rate of reaction, as there are more ions available to carry the current. This principle is crucial in industrial applications of electrolysis, where controlling the concentration of the electrolyte can optimize the yield and purity of the desired products.

Inert electrodes are used in electrolysis to provide a conductive surface for the electrochemical reactions without themselves participating in the reactions. This is crucial when the focus is on the products of electrolysis from the electrolyte, rather than the electrode material. Inert electrodes do not react or interfere with the electrolyte, ensuring that the products formed at the electrodes are solely due to the ions present in the electrolyte. Common materials used for inert electrodes include graphite and platinum. Graphite is a form of carbon that is conductive and relatively inexpensive, making it a popular choice for classroom experiments. Platinum, although more expensive, is highly resistant to corrosion and is used in industrial applications. The choice of inert electrode material can depend on factors like cost, availability, and the specific requirements of the electrolysis process, such as the need for high conductivity or resistance to chemical attack.