Electrolysis

The term electrolysis was first popularised in the 19th century by Michael Faraday. It is a process that helped in the study of chemical reactions in obtaining pure elements. Today, electrolysis is commercially important as it is used widely in separating or obtaining pure elements from naturally occurring sources such as ores.

Table of Contents

• What Is Electrolysis?
• Electrolytic Process
• Cell Potential or Voltage
• Faraday’s Law of Electrolysis
• Product of Electrolysis
• Factors Affecting Electrolysis
• Electrolysis Applications
• Electrolysis Problems with Solutions

What Is Electrolysis?

Electrolysis is defined as a process of decomposing ionic compounds into their elements by passing a direct electric current through the compound in a fluid form. The cations are reduced at the cathode, and anions are oxidized at the anode. The main components that are required to conduct electrolysis are an electrolyte, electrodes, and some form of external power source is also needed. Additionally, a partition, such as an ion-exchange membrane or a salt bridge, is also used, but this is optional. They are used mainly to keep the products from diffusing near the opposite electrode.

An acidified or salt-containing water can be decomposed by passing an electric current to its original elements, hydrogen and oxygen. Molten sodium chloride can be decomposed into sodium and chlorine atoms.

Electrolysis is usually done in a vessel named ‘electrolytic cell’ containing two electrodes (cathode and anode), connected to a direct current source and an electrolyte which is an ionic compound undergoing decomposition, in either molten form or in a dissolves state in a suitable solvent. Generally, electrodes that are made from metal, graphite and semiconductor materials are used. However, the choice of a suitable electrode is done based on the chemical reactivity between the electrode and electrolyte as well as the manufacturing cost.

Electrolytic Process

In the process of electrolysis, there is an interchange of ions and atoms due to the addition or removal of electrons from the external circuit. Basically, on passing current, cations move to the cathode, take electrons from the cathode (given by the supply source battery), and are discharged into the neutral atom. The neutral atom, if solid, is deposited on the cathode and, if gas, moves upwards. This is a reduction process, and the cation is reduced at the cathode.

Also Read: Water Electrolysis

At the same time, anions give up their extra electrons to the anode and are oxidized to neutral atoms at the anode. Electrons released by the anions travel across the electrical circuit and reach the cathode completing the circuit. Electrolysis involves a simultaneous oxidation reaction at the anode and a reduction reaction at the cathode.

For example, when an electric current is passed through molten sodium chloride, the sodium ion is attracted by the cathode, from which it takes an electrode and becomes a sodium atom.

Chloride ion reaches the anode, gives its electron, and becomes chlorine atoms to form chlorine molecules.

Na⁺(in electrolyte) + e^–(from cathode) → Na …. At Cathode

Cl^–(from electrolyte) → e^– + Cl → Cl₂ …. At Anode

The electrolysis process, while useful to get elemental forms from compounds directly, can also be used indirectly in the metallurgy of alkali and alkaline earth metals, purification of metals, deposition of metals, preparation of compounds, etc.

Cell Potential or Voltage

The minimum potential needed for the electrolysis process depends on the ability of the individual ions to absorb or release electrons. It is also sometimes described as decomposition potential or decomposition voltage which is the minimum voltage (difference in electrode potential) between the anode and cathode of an electrolytic cell that enables electrolysis to occur.

The voltage at which electrolysis is thermodynamically preferred is the difference of the electrode potentials, as calculated using the Nernst equation. Applying additional voltage, referred to as overpotential, can increase the rate of reaction and is often needed above the thermodynamic value. It is especially necessary for electrolysis reactions involving gases such as oxygen, hydrogen or chlorine.

This ability is measured as an electrode potential of the ions present in the electrolytic cell. The cell potential is the sum of the potential required for the reduction and oxidation reaction. The potential involved in various redox reactions is available in literature as standard reduction potential.

Reaction with positive redox cell potentials will only be feasible as per thermodynamic Gibbs free energy (or standard potential). Generally, the electrolysis is thermodynamically controlled.

In electrolysis, a potential equal to or slightly more than that is applied externally. The ions, which are stable and not reacting, are made to undergo reaction in the presence of externally applied potential. External potential, hence, causes an unfavourable reaction to take place. In electrolysis, chemical bonds connecting atoms are either made or broken; hence, electrolysis involves the conversion of electrical energy into chemical energy.

Faraday’s Law of Electrolysis

The amount of the redox reaction depends on the quantity of electricity flowing through the cell. The amount of reaction or the number of ions discharged is given by Faraday’s law of electrolysis. There are two laws: Faraday’s first law and second law.

Faraday’s first law can be summarised as follows:

Faraday’s second law compares the mass of different substances undergoing a change for the same current.

According to the second law,

Here, M and E are the changed mass and equivalent mass of the substances, respectively.

Product of Electrolysis

The products of electrolysis reactions depend on the oxidising and reducing species present in the electrolytic cell. Electrolysis will produce products present in the compound. When more than one cation and anions are present, each ion will compete for reduction and oxidations. Reactions with more positive redox potentials will be reduced or oxidized in preference to others.

Also Read: Nernst Equation

So, in spite of multiple redox couples present, only one can be reduced or oxidized. Sometimes, the ions that are reduced or oxidized may depend on their relative amount. In other words, the redox reaction and electrolysis may become kinetically controlled. In such cases, the product of analysis may differ on the relative concentration of the various ions present in the electrolyte.

For example, electrolysis of aqueous sodium chloride may give different products:

1. Hydrogen and chlorine
2. Hydrogen and oxygen
3. Hydrogen, oxygen and chlorine.

Factors Affecting Electrolysis

The factors that may affect electrolysis are listed below:

i) The nature of the electrode

ii) Nature and state of the electrolyte

iii) Nature and electrode potential of ions present in the electrolyte

iv) Overvoltage at the electrodes.

i) Nature and State of the Electrolyte

Electrolysis involves the movement of ions towards the oppositely charged electrodes. Naturally, the electrolyte should have mobile ions. In solids, ions are in specific positions and cannot move at ordinary temperatures. Hence, solids are unsuitable for electrolysis.

For electrolysis, electrolytes should be in the liquid form – molten or in solution with a suitable polar solvent. Sodium chloride will undergo electrolysis in the molten state or in an aqueous solution.

ii) Nature and Electrode Potential of Ions Present in the Electrolyte

• Electrolysis of electrolytes of two elemental ions is straightforward, giving the two elements on electrolysis. Molten sodium chloride gives sodium atoms and chlorine molecules.
• Electrolysis of radical ions does not give the elemental atoms.
• Electrolytes containing more than one ionic compound depend on the relative redox potentials.
• Electrolysis of aqueous solutions of electrolytes – Water molecules also can undergo redox reactions and will compete with redox reactions of the electrolyte ions.
• Electrolysis of molten sodium chloride gives sodium and chlorine. But, electrolysis of aqueous sodium chloride gives hydrogen and chlorine, and not sodium.

Also Check ⇒ Electrolytes

iii) Nature of the Electrode

For the same electrolyte, the nature of the electrolyte may give different products. When aqueous copper sulphate solution is, electrolyzed, the following redox reactions are possible.

At cathode: Reduction at pH =7

Cu²⁺ (aq) + 2e^– →Cu (s) E° = 0.34V and 2H₂O + 2e^–→H₂ + 2OH^– E° = -1.02V

At anode: Oxidation at pH = 7

Cu(s) →Cu²⁺ (aq) + 2e^– E° = – 0.34V and 2H₂O → O₂(g) + 4H⁺ + 4e^– E° = +1.4 V

At the cathode, out of the two electrodes, the reduction potential of copper ions is more positive than the reduction of water. So, irrespective of the electrode, copper ions from the electrolyte will be reduced and deposited on the cathode, increasing its mass. But, the reaction at the anode depends on the electrode.

Also Read: Electrochemical Cells

Electrolysis with inert electrodes like platinum, graphite, etc. – Inert electrodes do not react with the electrolyte or the products, so does not undergo any changes. Since the oxidation of water has more positive potential, oxygen will be evolved at the anode.

But, if the copper is used as an anode, it will react with the sulphate ion to retain the electrolyte concentration. So, there will not be any gas evolution; instead, the anode mass slowly decreases, going into the solution.

iv) Overvoltage at the Electrodes

The redox potential of electrolyte ions decides the electrolysis reactions and products. Sometimes, the redox potentials of some half-reactions during the electrolysis are more than the thermodynamic potentials. This excess voltage (over-voltage) of the half-reaction may make the reaction unfavourable and change the product of electrolysis.

In the hydrolysis of aqueous sodium chloride at the anode, two oxidation reactions can take place. The reduction potential of water and chloride is +0.82V and 0.1.36V, respectively.

2H₂O→O₂(g) + 4H⁺ + 4e^– E° = -0.82 V

2Cl^– → Cl₂ + 2e^– E = – 1.36V

Oxidation of water being more positive is more feasible, so the evolution of oxygen gas should happen at the anode. But, the evolution of oxygen from water has an overvoltage of -0.6V, making the voltage for the oxidation of water as -1.42V. Chloride oxidation is more positive than the net voltage of water oxidation. Chloride is oxidized to chlorine at the anode. Chlorine is liberated, and not oxygen, because of overvoltage.

Electrolysis Applications

Electrolysis, as stated above, is a process of converting the ions of a compound in a liquid state into their reduced or oxidized state by passing an electric current through the compound. Thus, electrolysis finds many applications, both in experimental and industrial products. Some of the important ones are given below:

1) Determination of equivalent weight of substances.
2) Metallurgy of alkali and alkaline earth metals.
3) Purification of metals.
4) Manufacture of pure gases.
5) Manufacture of compounds like sodium hydroxide, sodium carbonate, potassium chlorate etc.
6) Electroplating for corrosion resistance, ornaments etc.

We will discuss the different applications of electrolysis in detail below.

Determination of Equivalent Weight of Substances

We know Faraday’s second law states that the mass of substances deposited is proportional to their equivalent weight. The mass of any deposited substance can be, mathematically related as:

The equivalent mass of an unknown metal or substance can be calculated by passing a known current through the solutions and determining the mass of substances (M1 and M2) deposited in their respective cells. If the equivalent mass of one substance is known, the equivalent mass of the unknown substance can be calculated from the above equation.

Electrolysis of Molten Salts

Metallurgy of alkali and alkaline earth and third group metals ores of metal is concentrated and converted mostly to oxides. Oxides are reduced with reducing agents such as carbon, aluminium etc. Since alkali and alkali earth metals have the largest reduction potentials, any other metals or their compounds cannot reduce them.

The only way of isolation of alkali and alkali earth metals is to directly electrolyze their molten chlorides. Mixing with other halides, like calcium chlorides, reduces the melting point of pure halides.

Electrorefining – Purification of Metals

Metals obtained after concentration and reduction of ores have a purity of about 90-99%. An aqueous solution of the metal salt with the impure metal as the anode and the pure metal as the cathode is electrolyzed. The pure metal of more than 99% purity deposits on the cathode, and the impurities are collected at the bottom as mud. Copper and nickel are some examples of the metal purified by electrorefining.

Electroplating

An object can be coated to the required thickness with a select metal by electrolysis. The object to be coated is made of the cathode. An aqueous solution of the metal salt to be coated is the electrolyte. The same metal or any inert metal can be the anode. In electrolysis, metal ion from the electrolyte deposit on the object. The loss of metal ions in the solution will be compensated if the same metal made the anode.

The deposition can be used to protect the metal from corrosion for making ornaments, etc. Coating iron with metals like zinc, lead, chromium, and nickel improves the corrosion resistance of iron. Gold and silver coating on cheaper metals is used for making ornaments.

It is also used in electrochemical machining (ECM). Here, an electrolytic cathode is used as a shaped tool for removing material by anodic oxidation from a workpiece. ECM technique is often used for deburring or for putting a permanent mark or logo on metal surfaces like tools or knives.

Electro-forming

Electroforming is a process of making a replica of objects using electrolysis. The object to be replicated is pressed in wax to make a mould. Graphite powder is coated uniformly to make it conductive. This is used as a cathode, and the salt of the metal to be deposited is taken as the electrolyte. After getting the required coating by electrolysis, the wax and the graphite are melted away.

Manufacture of Pure Gases

Aqueous salts on hydrolysis yield different products depending on the relative concentrations of salt and water. Electrolysis of concentrated brine (sodium chloride) forms pure hydrogen and chlorine gases. Pure chlorine gas is collected in the Chlor-alkali industries by the electrolysis of brine aqueous solution.

Pure hydrogen and oxygen are obtained by hydrolysis of water in the presence of acid or base or inert salt of alkali and alkaline earth metals. The percentage of hydrogen for commercial use is manufactured by the electrolysis of water worldwide.

Continuous electrolysis of water removes all the normal hydrogen isotopes leaving the deuterium ions. The deuterium oxide leftover after electrolysis of normal water is ‘heavy water’. Heavy water is used as a moderator in nuclear reactors producing electrical energy from nuclear reactions.

Manufacture of Compounds

Compounds like sodium hydroxide, sodium hydrosulphite, potassium permanganate, potassium chlorate, ammonium per-sulphate, heavy water etc., are manufactured by electrolysis. Sodium hydroxide is a side product in the chloralkali industries, preparing chlorine gas by the electrolysis of brine.

Potassium permanganate is obtained by the electrolysis of potassium manganite solution. Ammonium sulphate or ammonium bisulphate on electrolysis forms ammonium persulphate.

Electrocrystallization

This is a specialised application of electrolysis. In this process, conductive crystals are grown on one of the electrodes from oxidized or reduced species that are generated in situ. This technique is popularly used to manufacture single crystals of low-dimensional electrical conductors, such as linear chain compounds or charge-transfer salts.

Electrolysis Problems with Solutions

1. An iron pipe with a 14 cm diameter and length of 1 metre is to be galvanized to a thickness of 0.01 cm using zinc nitrate solution and a current of 25 amp. What will be the loss of mass from the zinc anode, and what will be the time required for the electrolysis? The density of zinc is 7.14g/cm³. The equivalent weight of zinc is 32.8

Volume of the zinc coating = 2πrxl × 0.01 cm³

Mass of the zinc to be coated= V × d

The mass of the zinc lost from the anode = 440g.

By Faraday’s first law of electrolysis,

= 51772sec

= 14.4hrs

2. A current of 3 amperes is passed through neutral water containing a small amount of sodium sulphate for one hour. Calculate the amount of hydrogen liberated at one atmosphere.

Number of coulombs passed = 3 × 60 × 60

Half reactions of the hydrolysis of water are –

At cathode: 2H₂O + 2e^– → H₂(g) + 2OH^–

At anode: 2H₂O → O₂(g) + 4H⁺ + 4e^–

The net reaction of electrolysis of water is, 2H₂O → 2H₂(g) + O₂(g)

Four faradays or 4 x 96485 coulomb of electric current liberates 2 moles or 2× 22.4L of hydrogen gas.

10800 coulombs of electric current liberates hydrogen

Total pressure above water = pressure of water vapour + pressure of hydrogen

The pressure of hydrogen = Total pressure above water – pressure of water vapour

The pressure of hydrogen = 1atm – 0.0316atm =0.9684atm

So, volume of hydrogen liberated 

Electrochemistry – Important Topics

Electrochemistry – Important Questions