Electrolysis: what it is and its application

Electrolysis is a process that uses direct electric current to drive chemical reactions that don’t occur spontaneously.
The foundations of electrolysis were laid in the late 18th century when scientists began experimenting with electricity and its effects on chemical substances.
Remarkably, Nicholson and Carlisle conducted groundbreaking experiments demonstrating water electrolysis using a voltaic pile, effectively converting water into hydrogen and oxygen gases.
This experiment provided the first evidence of a chemical transformation initiated through an electrical current.

Nearly a century later in 1883, Michael Faraday found a quantitative relationship between electricity and mass.
Faraday's laws of electrolysis (1), formulated based on his extensive experimental work, provided a quantitative understanding of the relationship between the amount of substance produced during electrolysis and the quantity of the electricity supplied:

in equation (1), m is the mass-produced, I is the current intensity, Δt is the time frame during which the electric current is applied, Z is the ion charge, and F is the Faraday constant.

Experimentally, electrolysis occurs in an electrochemical cell made of two electrodes – the anode and the cathode – separated by a diaphragm or a membrane.
Either the anode and the cathode are immersed in a solution of positively and negatively charged particles called ions. The solution of ions is the so-called electrolyte, and more specifically, the solution in which the cathode is immersed is called the catholyte, whereas the anodic one is called the anolyte.

During electrolysis, when an electric current is applied across the electrodes, ions within the electrolyte migrate toward the electrodes under the influence of the electric field. At the electrodes, they undergo oxidation or reduction reactions, depending on their charge and the electrode's polarity.

Specifically, the electrode at which oxidation occurs is termed the anode and it is positively charged, while the electrode at which reduction occurs is called the cathode and it’s negatively charged.
The oxidation reaction means that chemical substances lose electrons (2), whereas reduction means that a chemical species gains electrons (3):

De Nora has an extensive knowledge of electrochemistry and electrolysis gained through its centenarian history and has developed several technologies in this field, especially in the chloro-alkali, electrowinning of non-ferrous metals, and water electrolysis markets.

The chlor-alkali process is an industrial process for the electrolysis of sodium chloride (NaCl) to produce chlorine and sodium hydroxide (caustic soda). In this process, saturated brine is supplied at the anode of an electrochemical cell where the chloride ions are oxidized, losing electrons to become chlorine gas (4):

At the cathode, water is reduced by the electrons provided by the cathode
to hydrogen gas, releasing hydroxyl ions into the solution (5):

An ion-exchange membrane is placed at the center of the electrochemical cell.
Its role is to separate the anolyte from the catholyte and shuttle Na+ ions from the anode to the cathode, thus originating the NaOH. The overall chemical reaction is displayed below (6).

The electrowinning of non-ferrous metals is another application of the electrolysis process. In fact, metal ores contain metals of interest (cobalt, copper, nickel) in their oxidized states.
Therefore, the purpose of metallurgic operations is to chemically reduce and extract them into their pure forms. The electrowinning process begins with the dissolution of the ore with chloric or sulphuric acid and, once purified, using these solutions as electrolytes of an electrochemical cell. Immersed in the solution, we have an anode, while a plate of pure metal is used as the cathode.
For example, in the electrowinning of nickel, we have an anode, and a plate of pure nickel is used as the cathode. They are both immersed in a solution of nickel chloride or sulphate (obtained by dissolution of the ore with chloric or sulphuric acid, respectively). When an electric current is applied across the electrodes, the nickel ions migrate towards and are reduced at the cathode to form the pure metal (7):

The reaction at the anode is the formation of chlorine gas, as displayed in equation (4) if the electrolyte is made of nickel chloride, whereas we have an oxygen evolution reaction (8) when a nickel sulfate solution is used as electrolyte.

Water electrolysis to produce green hydrogen is another application of electrolysis and consists of breaking down water molecules into their constituents – hydrogen and oxygen – by means of renewable sources as an energy supply (9). Leveraging on the centenarian knowledge gained through the chloro-alkali and electrowinning processes, De Nora broke into the green technology market with novel solutions to perform water electrolysis at scale.

In conclusion, electrolysis is the process by which an electric current initiates chemical reactions. These chemical reactions occur within an electrochemical cell constituted by an anode and a cathode separated by a diaphragm or a membrane.
⁠Electrolysis has a wide range of industrial applications, including producing chlorine gas and sodium hydroxide (chlor-alkali process), electrowinning of non-ferrous metals, and the production of green hydrogen through water electrolysis.