Electrolysis of Water: Diagram, Experiment & Chemical Equation

Electrolysis of Water: With the ever-increasing need for fossil fuels, scientists are always looking for new ways to generate energy. Thanks to hydrogen, which does not pollute the air when burned. Finding cost-effective ways to create hydrogen has been a challenge. However, its generation by solar cells splitting water molecules promises to be a potential source of power in the future. In this article, we’ll learn more about this procedure.

 Electrolysis is currently used in all the prime industries. Around 5% of the hydrogen gas is produced from the process electrolysis. Nowadays, natural gas is used in the process of electrolysis. Even in space stations, water electrolysis is used to produce oxygen. Check out the article to know more about the electrolysis of water, affecting factors, Electrolysis of Pure Water and many more.

Electrolysis of Water: Overview

The electrolysis of water is carried out in an electrolytic cell consisting of a pair of platinum electrodes immersed in water. To this water, a small amount of an electrolyte such as \({{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4}\) has been added. The addition of an electrolyte is necessary because pure water lacks ions responsible for conducting electricity. 

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On passing an electric current through the solution, water decomposes to oxygen gas and hydrogen ions at the anode. In contrast, at the cathode, water is reduced to hydrogen gas and hydroxide ions.

At anode(Oxidation): \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l)}} \to {{\rm{O}}_{\rm{2}}}{\rm{(g) + 4}}{{\rm{H}}^{\rm{ + }}}{\rm{(aq) + 4}}{{\rm{e}}^ – }\quad {{\rm{E}}^{\rm{0}}}{\rm{ =\,  – 1}}{\rm{.23\;V}}\)

At cathode(Reduction): \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l) + 2}}{{\rm{e}}^{\rm{ – }}} \to {{\rm{H}}_{\rm{2}}}{\rm{(g) + 2O}}{{\rm{H}}^{\rm{ – }}}{\rm{(aq)}}\quad {{\rm{E}}^{\rm{0}}}{\rm{ =\,  – 0}}{\rm{.83\;V}}\)

Overall reaction: \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l)}} \to {{\rm{O}}_{\rm{2}}}{\rm{(g) + 2}}{{\rm{H}}_{\rm{2}}}{\rm{(g)}}\quad {\rm{E}}_{{\rm{cell }}}^{\rm{0}}{\rm{ =\,  – 2}}{\rm{.06\;V}}\)

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In order to obtain the overall reaction, the reduction half-reaction was multiplied by two to equalise the electrons. The hydrogen ion and hydroxide ions produced in two half-reactions combine to form water. The \({{\rm{H}}_2}{\rm{S}}{{\rm{O}}_4}\) is not consumed in the reaction.

Hence, the electrolysis of water produces hydrogen gas at the cathode and oxygen gas at the anode. 

However, the electrolysis of water is not simple and easy for many reasons.

(i) Pure water is a bad conductor of electricity. Hence, it is very weakly dissociated into hydrogen ions and hydroxide ions. So, electrolysis of pure water will be a prolonged process.

(ii) The hydrogen ion produced at anode associates with another water molecule to form hydronium ions. So, any hydroxide ions moving towards the anode will be neutralised by the hydronium ions, even before it reaches the anode to form oxygen gas. Similarly, hydroxyl ions neutralise hydrogen ions present near the cathode and will not be reduced to hydrogen. Hence, the electrolysis of water to hydrogen and oxygen will be very small. 

Factors Affecting the Efficiency of Electrolysis

The efficiency of electrolysis or the electron transfer depends on many factors such as;

i) The number of cations and anions present in the solution.

ii) Mobility rate of the ions to reach the electrode.

iii) Activation energy needed for the electron transfer from the electrode to the electrolyte ions.

iv) Nature of electrode and electrolyte.

v) standard electrode potentials of the different oxidising and reducing agents present in the electrolytic cell.

Practice Exam Questions

Electrolysis of Pure Water (pH =7)

Half reactions in the electrolysis of pure water at \({\rm{pH}} = 7,\), and at \({25^ \circ }\) are-

At cathode(Reduction): \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l) + 2}}{{\rm{e}}^{\rm{ – }}} \to {{\rm{H}}_{\rm{2}}}{\rm{(g) + 2O}}{{\rm{H}}^{\rm{ – }}}{\rm{(aq)}}\quad {{\rm{E}}^{\rm{0}}}{\rm{ =\,  – 0}}{\rm{.83\;V}}\)

At anode(Oxidation): \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l)}} \to {{\rm{O}}_{\rm{2}}}{\rm{(\;g) + 4}}{{\rm{H}}^{\rm{ + }}}{\rm{(aq) + 4}}{{\rm{e}}^{\rm{ – }}}\quad {{\rm{E}}^{\rm{0}}}{\rm{ =\,  – 1}}{\rm{.23\;V}}\)

Overall reaction: \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l)}} \to {{\rm{O}}_{\rm{2}}}{\rm{(g) + 2}}{{\rm{H}}_{\rm{2}}}{\rm{(g)}}\quad {\rm{E}}_{{\rm{cell }}}^{\rm{0}}{\rm{ =\,  – 2}}{\rm{.06\;V}}\)

The cell potential of electrolysis of pure water is negative and hence is thermodynamically unfavourable. Because of the low concentration of ions an extra voltage (Overvoltage) of about \(0.4 – 0.6\;{\rm{V}}\) at each electrode is needed for the electrons to cross the interface.

In practice, continuous electrolysis of pure water is possible only at an external voltage of \(2.4\;{\rm{V}}\)

Water Electrolysis in the Presence of Acids (pH lower than 7)

Acids dissociate to give additional hydrogen ions that are reduced at the cathode, while water will be oxidised at the anode. Half reactions in an acidic medium are:

At cathode(Reduction): \({\rm{2}}{{\rm{H}}^{\rm{ + }}}{\rm{ + 2}}{{\rm{e}}^{\rm{ – }}} \to {{\rm{H}}_{\rm{2}}}{\rm{(g)}}\quad {{\rm{E}}^{\rm{0}}}{\rm{ = + 0}}{\rm{.0V}}\)

At anode(Oxidation): \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l)}} \to {{\rm{O}}_{\rm{2}}}{\rm{(g) + 4}}{{\rm{H}}^{\rm{ + }}}{\rm{(aq) + 4}}{{\rm{e}}^{\rm{ – }}}\quad {{\rm{E}}^{\rm{0}}}{\rm{ =\,  – 1}}{\rm{.23\;V}}\)

Overall reaction: \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l)}} \to {{\rm{O}}_{\rm{2}}}{\rm{(g) + 2}}{{\rm{H}}_{\rm{2}}}{\rm{(g)}}\quad {\rm{E}}_{{\rm{cell}}}^{\rm{0}}{\rm{ =\,  – 1}}{\rm{.23\;V}}\)

The electrolysis takes place at a much lower potential than pure water.

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Water Electrolysis in the Presence of a Base (pH higher than 7)

Bases dissociate to give additional hydroxyl ions that release their electrons at the anode to oxidise into oxygen gas, while electrons at cathode reduce water molecules into hydrogen gas. Half reactions of electrolysis in the presence of a base are-

At cathode(Reduction): \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l) + 2}}{{\rm{e}}^{\rm{ – }}} \to {{\rm{H}}_{\rm{2}}}{\rm{(\;g) + 2O}}{{\rm{H}}^{\rm{ – }}}\quad {{\rm{E}}^{\rm{0}}}{\rm{ =\,  – 0}}{\rm{.83\;V}}\)

At anode(Oxidation): \({\rm{4O}}{{\rm{H}}^{\rm{ – }}} \to {{\rm{O}}_{\rm{2}}}{\rm{(g) + 2}}{{\rm{H}}_{\rm{2}}}{\rm{O + 4}}{{\rm{e}}^{\rm{ – }}}\quad {{\rm{E}}^{\rm{0}}}{\rm{ =\,  – 0}}{\rm{.4V}}\)

Overall reaction: \({\rm{2}}{{\rm{H}}_{\rm{2}}}{\rm{O(l)}} \to {{\rm{O}}_{\rm{2}}}{\rm{(g) + 2}}{{\rm{H}}_{\rm{2}}}{\rm{(g)}}\quad {\rm{E}}_{{\rm{cell }}}^{\rm{0}}{\rm{ =\,  – 1}}{\rm{.23\;V}}\)

Like electrolysis in an acid medium, electrolysis in the basic medium also needs much lower potential.

Applications of Electrolysis of Water

1. Electrolysis of water is used to generate oxygen for the International Space Station.

2. Hydrogen produced as a byproduct in the chlor alkali process is used to produce speciality chemicals or various other small-scale applications.

3. Hydrogen gas released in this way can be used as hydrogen fuel or remixed with oxygen to create oxyhydrogen gas used in welding and other applications.

4.Hydrogen produced through electrolysis of water is used to manufacture heavy water.

5. About five per cent of hydrogen gas produced worldwide is created by electrolysis. 

Summary

Hydrogen is an essential alternative to fossil fuels. With the booming population, the need for fuels is exponentially increasing. Hence, it is important to synthesise it on a large scale without any damage to the environment. We can synthesise hydrogen through the electrolysis of water. In this article, we learnt the process of electrolysis and factors that influence its rate. We also learnt some of its applications. 

Frequently Asked Questions (FAQs) on Electrolysis of Water

The most commonly asked questions on Electrolysis of water are answered here:

Q.1: What happens during the electrolysis of water?
Ans: During the electrolysis of water, decomposition of water into oxygen gas and hydrogen gas takes place. This happens by passing an electric current through the solution. Hydrogen gas is produced at the cathode, whereas oxygen gas is produced at the anode. 

Q.2: What are the observations of electrolysis of water?
Ans: The gases collected during the electrolysis of water can be tested with a glowing splint (which will flame up in the presence of oxygen) and a burning splint (which will ignite hydrogen with a pop sound). \({{\rm{H}}_2}\) gas is produced at twice the rate of \({{\rm{O}}_2}\) gas in this reaction.

Q.3: Which type of chemical reaction is the electrolysis of water?
Ans: Electrolysis is an example of a decomposition reaction. This is because water splits into its constituent elements, hydrogen and oxygen, under the effect of an electric current.

Q.4: Which electrodes are used in the electrolysis of water?
Ans: Platinum electrodes are used in the electrolysis of water.

Q.5: What is the conclusion of the electrolysis of water?
Ans: The conclusion of the electrolysis of water is that hydrogen gas is produced at the cathode, whereas oxygen gas is produced at the anode.

Q.6: Does electrolysis destroy water?
Ans: Yes, electrolysis breaks down water into its constituting ions \({{\rm{H}}^{\rm{ + }}}\) and \({\rm{O}}{{\rm{H}}^{\rm{ – }}}\). These ions can join up to form water molecules, but they are not allowed to do so in the electrolysis process.

Q.7: What voltage is best for the electrolysis of water in the acidic medium?
Ans: The minimum voltage necessary for electrolysis of water in presence of acid is about \({\rm{1}}{\rm{.23\;V}}\) If electrolysis is carried out at a high temperature, this voltage reduces.

Study Faraday’s Laws Of Electrolysis

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