What is Electrolysis?

Electrolysis is defined as the process of breaking down ionic compounds into their constituent elements by passing a direct electric current through the compound in a fluid state. The cathode reduces cations, while the anode oxidizes anions. An electrolyte, electrodes, and some form of external power source are the main components required for conducting electrolysis. A partition, such as an ion-exchange membrane or a salt bridge, may also be used, but this is optional. These are primarily used to prevent the products from diffusing near the opposing electrode.

Where Does Electrolysis Occur?

The battery of your phone or laptop is running down. What do you do? You plug it in to recharge.

Ever wondered what is actually happening when you put these devices on charge? 
Correct, electrolysis occurs. It is a process which uses electric current to dissociate compounds and carry out reactions that are otherwise non-spontaneous. Batteries utilize a chemical reaction and provide electrical energy. When the reactants are used, batteries start to slow down and then we recharge the battery reversing the reaction. While recharging, the battery is an electrolytic cell where electrolysis is carrying out a non-spontaneous reaction to replenish back the reactants. Not only this, the silver or gold-plated jewellery, copper-plated utensils, and myriad of things around you, like chromium-plated car bumpers, nickel- or chromium-plated water taps, tin plated iron cans, etc. are plated with metals to avoid them from tarnishing, corrosion or just for aesthetic purposes, all are a result of electrolysis. So, let’s learn about electrolysis and try to understand some hidden chemistry in our surroundings related to it.              

The term electrolysis was coined by Michael Faraday in the 19^th century. Electro signifies electricity and lysis from the Greek origin means ‘to split’. Refer to the figure (Fig 1) below for the illustration of an electrolytic cell that is used to perform electrolysis.

Parts of an Electrolytic Cell

Electrodes

There are two electrodes connected to the ends of a battery (DC source). One electrode, connected to the negative end of the battery is negative and is called cathode. The other end connected to the positive terminal of the battery is positive and is called anode. Electrodes conduct electrons in and out of the cell and also provide the surface for reactions that occur during electrolysis. Electrodes may be made up of materials, such as zinc, copper, and silver. These electrodes are called active electrodes as they participate in the electrolytic reaction. Another type of electrodes is inert electrodes that do not participate in the reaction and are made up of materials like graphite and platinum.

Electrolyte

Within the cell, the substance whose electrolysis is carried out serves as the medium for the conduction of current between the electrodes. These substances have freely moving ions in their molten or aqueous state, thus conducting electricity. These substances are called electrolytes. An electrolyte can be defined as a substance that furnishes cations (positively charged ions) and anions (negatively charged ions) in molten state or in aqueous solutions. For example, sodium chloride furnishes sodium ions (cations) and chloride ions (anions).

Working of an Electrolytic Cell

Now that we know what an electrolytic cell is made of, let’s see how these parts work together (Fig 1). The underlying basic concept of electrolysis is the interchange of atoms and ions by taking away or providing the electrons from external circuit. The cell is filled with the electrolyte and a potential difference sufficient to carry out the electrolysis is applied. Electrons are supplied to the electrode connected to the negative terminal of the battery. This electrode becomes negatively charged (cathode). Electrons are also pulled out of the electrode connected to the positive end of the battery making it positively charged (anode). Within the solution, electrolyte ions carry the current. The positive cations move to the negative cathode, where they accept available electrons and neutralize on the cathode. Negatively charged anions from the electrolyte move towards the anode where they lose electrons and neutralize.  

Let us take a hypothetical example of an electrolyte MN, which ionizes as follows:

$MN \to M^{x+} + N^{x-}$

On applying the electric potential, the cation M^x+ moves towards the cathode (negative) and gains available electrons to form M, i.e., it undergoes reduction (gaining of electrons). The anion N^x- moves towards the anode and loses electrons to form N, i.e., it undergoes oxidation (losing electrons). We can write the electrode reactions as:

Reduction half reaction at cathode:

 $M^{x+} + x{e}^{-} \to M$  

Oxidation half reaction at anode:

$N^{x-} \to N + x{e}^{-}$        

The overall reactions is:

$M^{x+} + N^{x-} \to M + N$ 

Thus, electrolysis carries out a redox reaction, where reduction half reaction takes place at cathode, and oxidation half reaction occurs at anode. The overall reaction which otherwise would have been non-spontaneous is driven by the current.

Electrolysis of Molten Sodium Chloride

Electrolyte: Molten sodium chloride

Electrodes: Platinum electrodes (inert)

Freely moving electrolyte ions: Na⁺ and Cl^-

Working of the cell: When molten NaCl is subjected to electrolysis, sodium ions (Na⁺) move towards cathode, gain electrons and form sodium (Na). The chloride ions (Cl^-) move towards anode, lose electrons and form Cl₂ gas that evolves at the electrode (Fig 2).

Ionization of NaCl:

$NaCl \to N{a}^{+} + C{l}^{-}$

Reduction half reaction at cathode:

$N{a}^{+} + e^{-} \to Na$

Oxidation half reaction at anode:

$C{l}^{-} \to \frac{1}{2}C{l}_2 + e^{-}$

The overall reaction:

$N{a}^{+} + C{l}^{-} \to Na + \frac{1}{2}C{l}_2$

Electrolysis of Water (H₂O)

Electrolysis of water is the breaking down of water into hydrogen and oxygen. This technique is useful as hydrogen and oxygen can be further used. Oxygen that astronauts or those who travel on submarines need to breathe, can get through this way! Hydrogen so produced is a clean (carbon free production) and renewable source that can be used as a fuel.

It must be noted that water is very feebly dissociated into its hydrogen and hydroxide ions. It is a poor conductor of electricity. Electrolysis involves electrolyte ions for current flow, and since the amount of ions produced by water is small, electrolysis of pure water is a very slow process. We need to increase the conductivity of water for carrying out its electrolysis efficiently and that is done by adding an acid, base or salt to the water. These substances dissociate into their ions in water increasing its conductivity.    

For electrolysis of water, two inert electrodes, like that made up of platinum are placed in the electrolytic cell containing water and acid/base/salt (Fig 3).

Water Electrolysis with Acid

H⁺ ions from acid move towards cathode and get reduced to form hydrogen gas, while water gets oxidized at anode to form oxygen gas. Acids like H₂SO₄ and HNO₃ are used. Sulphate (SO₄^2-) and nitrate (NO₃^-) ions have electrode reduction potential higher than that of H₂O, so they do not interfere with production of oxygen gas at anode.

Reduction half reaction at cathode:

 $2H^{+} + e^{-} \to H_2\left(g\right)$

Oxidation half reaction at anode:

$2H_{2}O \to O_2\left(g\right) + 4H^{+} + 4e^{-}$

Overall reaction:

$2H_{2}O \to O_2\left(g\right) + 2H_2\left(g\right)$

Water Electrolysis with Base

OH^- ions from the base move towards anode, lose electrons and get oxidized to form oxygen gas while water at cathode gets reduced by accepting electrons to form hydrogen gas. Bases like NaOH and KOH are used. Sodium (Na⁺) and potassium (K⁺) ions have electrode reduction potential lesser than H⁺ ions, so they do not compete with H₂O at cathode.

Reduction half reaction at cathode:

$2H_{2}O + 2e^{-} \to H_2\left(g\right)$

Oxidation half reaction at anode:

$4O{H}^{-} \to O_2 +2H_{2}O + 4e^{-}$

Overall reaction:

$2H_{2}O \to O_2\left(g\right) + 2H_2\left(g\right)$

Water Electrolysis with Salt

Cation and anion generated from the salt may also attract towards the electrodes. So, the selection of salt is also important. Salts like NaNO₃ can be used. Neither Na⁺ nor NO₃^- ions interfere with the production of H₂ and O₂ gas.

Reduction half reaction at cathode:

$2H_{2}O + 2e^{-} \to H_2\left(g\right)$

Oxidation half reaction at cathode:

$2H_{2}O \to O_2\left(g\right) + 4H^{+} + 4e^{-}$

Overall reaction:

$2H_{2}O \to O_2\left(g\right) + 2H_2\left(g\right)$

At industrial level, hydrogen gas is produced by electrolysis of water using electrolyzers. It is an unit that contains a cathode, an anode and a membrane. There are different types of electrolyzers, like polymer electrolyte membrane electrolyzer, alkaline electrolyzer, which function in different ways.

Body Hair Removal Using Electrolysis

The practice of removing human body hair permanently using electric current by electrolysis is called electrology. A qualified professional called electrologist inserts a fine probe into each hair follicle. Current is passed through the probe to the follicle. Electrolysis of salt and water (aqueous NaCl) is carried out producing sodium hydroxide. Sodium hydroxide kills the cells required for hair growth. Hair removal by electrolysis is FDA-approved.

Electroplating

Coating one metal over another using electric current is called electroplating. It helps prevent corrosion, increases life span of the metal beneath, and provides insulation. It serves a good aesthetic purpose at a lower price, as costly metal can be coated on an economical object surface. Objects are electroplated with chromium, nickel, gold or silver.    

In an electrolytic cell used to carry out electroplating, anode is the metal that has to be coated. Cathode is the object that has to be coated. The electrolyte is the solution of the coating metal. See Fig 4 for copper plating. Cu loses electrons on anode, and migrates to cathode as Cu²⁺ ions where it gets deposited as Cu by accepting electrons.

In a similar way, impure metals can be purified where the metal block is taken as anode and a pure metal strip is taken as cathode. Electrolysis is also used in extracting metals from their naturally occurring ores. Examples of such metals are potassium, sodium, calcium, magnesium and aluminium.

Context and Applications

In Industry you can see wide applications of the process of electrolysis. Some of these applications include manufacturing of heavy water, production of hydrogen, electroplating, metal refining.

This Crystal lattices and unit cells are studied under crystallography. Knowledge of unit cells helps in accurately evaluating the structure of a solid. You will study these in your higher studies in Chemistry. This topic is significant in the professional exams for both undergraduate and graduate courses, especially for

Bachelors of Science in chemistry

Master of Science in chemistry

Practice Problems

Q1. Which of the following is the preferred electrode when it does not participate in the chemical reaction?

1. Zinc
2. Iron
3. Copper
4. Graphite

Answer: Graphite

Q2. Which of the following statements is NOT correct about electrolytic cells?

1. At cathode, a reduction half reaction takes place.
2. At anode, a oxidation half reaction takes place.
3. The anion undergoes oxidation at cathode.
4. The cation undergoes reduction at the cathode.

Answer: The anion undergoes oxidation at cathode.

Q3. Which product is formed at the anode in the electrolysis of molten NaCl?

1. Oxygen gas
2. Sodium metal
3. Hydrogen gas
4. Chlorine gas

Answer: Chlorine gas

Q4. Which of the following half-reaction occurs at the cathode during the electrolysis of water with base?

1. 2H₂O + 2e^–→ H₂(g)
2. 2H⁺ + e^–→ H₂ (g)
3. 4OH^–→ O₂ + 2H₂O + 4e^–
4. 2H₂O + 2e^–→ H₂(g)

Answer: 2H₂O + 2e^–→ H₂(g) 

Q5. In electroplating, the metal is dissolved from the _____________and plated on the _____________.

1. cathode, anode
2. anode, cathode
3. anode, anode
4. cathode, cathode

Answer: anode, cathode