Chemistry Learner

It's all about chemistry.

  • Chemical Bonds
  • Chemical Reactions
  • Materials Chemistry
  • Organic Chemistry
  • Periodic Trends
  • Periodic Table Groups
  • How to Read Periodic Table
  • Naming Covalent Compounds Worksheets
  • Net Ionic Equation Worksheets
  • Types of Chemical Reactions Worksheets
  • Word Equations Worksheets
  • Valence Electrons Worksheets
  • Graphing Periodic Trends Worksheets
  • Periodic Trends Ionization Energy Worksheets
  • Atomic Structure And Isotopes Worksheets

Electrolysis

What is electrolysis, how does electrolysis work, examples of electrolysis [1-4], laws of electrolysis [1], applications of electrolysis [6].

Electrolysis separates chemically bonded ionic substances and compounds by passing an electric current through them. It uses a direct current (DC) to drive a non-spontaneous reaction that occurs during the process [1-4] .

English physicist Michael Faraday popularized electrolysis in the 19 th century.

Electrolysis is carried out in an electrolytic cell consisting of a positively charged anode and a negatively charged cathode. For electrolysis to occur, the compound must contain ions. The presence of an electrolyte is essential since it is made up of ionic compounds that have free ions. An external power source like a battery is also required to power the carbon electrodes and drive the cell. The electrical energy is then converted into chemical energy, and the cell does electrical work [1-4] .

Ion Interchange

The primary mechanism of electrolysis is the interchange of ions and atoms, resulting in a redox reaction. A chemical change takes place when atoms and ions lose or gain electrons. Positively charged ions or cations move towards the negatively charged cathode. They accept electrons and neutralize. This process is known as reduction, and the cations are said to be reduced.

On the other hand, negatively charged ions move towards the positively charged anode. They lose electrons and get oxidized, a process known as oxidation. The formation of neutral atoms from ions at the electrodes is known as discharging.

Ionic Half Equations

Ionic half-equations represent the process of oxidation and reduction. Suppose M + is a metal ion that gains electrons (e – ) at the cathode to form a neutral atom (M). Then, the half-reaction is given by,

M + + e – → M

Suppose X – represents a negatively charged nonmetal that loses electrons at the anode and transforms to a neutral atom (X). Then, the half-equation is given by,

X – → X + e –

write 3 application of electrolysis

Predicting Products of Electrolysis

At the electrodes, ions gain or lose electrons, become neutral, and separate from the solution. It is generally easy to predict the products during the electrolysis of molten electrolytes because the compounds split into their elements. The metal part gets deposited as a solid metal. The nonmetal part is released as gas. The electrolysis of aqueous ionic compounds is complicated. Water produces hydrogen ions (H + ) and hydroxide ions (OH – ) and competes with the primary compound. Hence, hydrogen (H 2 ) and oxygen (O 2 ) gasses can also liberate [5] .

1. Water (H 2 O)

Water can undergo electrolysis in the presence of an electrolyte like acid or base. The presence of acid improves the electrical conductivity by increasing the hydrogen ion (H + ) concentration. Examples of such electrolytes are sulfuric acid (H 2 SO 4 ) and salt of sodium nitrate (NaNO 3 ). The half-reactions are given below.

At cathode:

2 H + (aq.) + 2 e – → H 2 (g)

Hydrogen gas (H 2 ) will be liberated at the cathode.

2 H 2 O (l) → O 2 (g) + 4 H + (aq.) + 4 e –

Oxygen gas (O 2 ) will be liberated at the anode.

The overall chemical reaction is,

2 H 2 O (l) → 2 H 2 (g) + O 2 (g)

For two moles of water, two moles of hydrogen and one mole of oxygen are liberated. The number of moles of hydrogen generated is twice the oxygen. Also, charges are transferred between the electrodes and the electrolyte. For every mole of hydrogen, 2 electrons are transferred from the cathode to the electrolyte. For every mole of oxygen, 4 electrons are transferred from the electrolyte to the anode.

2. Aqueous Sodium Chloride (NaCl)

For electrolysis of aqueous sodium chloride solution, one has to factor in the electrolysis of water. Since water can both oxidize and reduce, it will compete with sodium chloride. However, since sodium is more reactive than hydrogen, hydrogen gas will be released at the cathode.

2 H 2 O (l) + 2 e – → H 2 (g) + 2 OH – (aq.)

2 Cl – (aq.) → Cl 2 (g) + 2 e –

The overall chemical reaction can be written as,

2 H 2 O (l) + 2 Cl – (aq.) → H 2 (g) + Cl 2 (g) + 2 OH – (aq.)

Therefore, the electrolysis of an aqueous sodium chloride solution yields sodium hydroxide, hydrogen, and chlorine.

3. Molten Sodium Chloride (NaCl)

Sodium chloride is melted to its molten state above 800 ˚C before performing electrolysis.

Positively charged sodium ions (Na + ) migrate towards the negatively charged cathode and reduce to sodium atoms (Na), forming sodium metal.

Na + (l) + e – → Na (l)

Negatively charged chloride ions (Cl – ) migrate the other way towards the anode and form chlorine gas (Cl 2 ).

2 Cl – (l) → Cl 2 (g) + 2e –

The overall chemical reaction is given by,

2 NaCl (l) → 2 Na (s) + Cl 2 (g)

In summary, the electrolysis of molten sodium chloride produces metallic sodium and chlorine gas.

During electrolysis, charges transfer between the electrodes and the electrolyte and flow through the electrolyte. The amount of charge flowing per unit time is the current. Based on his experiments, Faraday proposed two laws governing the current and the weight of products formed at the electrodes.

  • The weight of products formed during electrolysis is proportional to the current sent through the electrolyte.
  • If the current is kept fixed, the weight of the products formed is proportional to each product’s equivalent weight.

The equivalent weight of each product is defined as the molar mass divided by the number of electrons required to neutralize it.

1. Extracting, purifying, or cleaning metals : The anode is the impure metal or ore, and the electrolyte is the salt of the metal to be extracted. The pure metal is deposited at the cathode. Aluminum is extracted in this way. Copper can be purified by taking a dilute aqueous solution of copper sulfate and sulfuric acid as the electrolyte.

2. Production of pure chemicals : Many chemicals like sodium hydroxide, caustic soda, potassium permanganate, and potassium chlorate, as well as gasses like oxygen and chlorine, are produced in the industry.

3. Electroplating : It is the process of ornamenting cheap metals like iron with precious metals such nickel, chromium, gold, or silver. It is used for making jewelry, utensils, and auto parts.

4. Electroforming : It reproduces objects like coins and medals by electrodepositing over a mold.

5. Determination of equivalent mass : From the equivalent mass of one metal, one can calculate the equivalent mass of another metal from electrolysis laws.

6. Rust removal: It is used for metal restoration and can remove rust from iron.

  • Courses.lumenlearning.com
  • Opentextbc.ca
  • Chemistrytalk.org
  • Chemed.chem.purdue.edu
  • Tutorhelpdesk.com

Related Articles

Mass Spectrometry

Mass Spectrometry

Gravimetric Analysis

Gravimetric Analysis

write 3 application of electrolysis

Galvanization

write 3 application of electrolysis

Electrochemical Series

6 responses to “Electrolysis”

' src=

Perfect 👍, just what I needed

' src=

wow! very understandable

' src=

Perfect explanation!

' src=

Good. Just what I needed. Thank you.

Leave a Reply Cancel reply

Your email address will not be published. Required fields are marked *

Trending Topics

© 2024 ( Chemistry Learner )

Chemistry Steps

Chemistry Steps

The Electrolysis of Molten NaCl

General Chemistry

Electrochemistry.

We have seen earlier that galvanic cells have a positive cell potential (negative Gibbs free energy) which makes them spontaneous . This means that, just like for any other process, the reverse reaction of any spontaneous redox reaction will be nonspontaneous as it will have a negative cell potential and a positive Gibbs free energy .

For example, the reaction between sodium and chlorine is violent as it is very exothermic and therefore, we can also predict that the overall cell potential will be a positive value.

write 3 application of electrolysis

The cell potential is estimated based on the half-reactions under standard conditions.

We can explain this reaction simply based on the fact that Na is a metal and likes to lose an electron to become a cation which is easily achieved when there is a nonmetal with great electronegativity such as Cl. So, in a simple scheme, we can visualize this as Cl pulling the electron from Na because it likes electrons a lot more than Na does.

Now, here is the question: if the reverse reaction is nonspontaneous, (how) can we force it to occur? In the case of Na and Cl 2 , the target reverse reaction would be:

write 3 application of electrolysis

Can we make Cl – ions give the electrons back to Na + and prepare Na metal and Cl 2 gas?

Remember, from thermodynamics that if a process is endothermic (more accurately endergonic if we are talking of the Gibbs free energy ), then we need to supply the necessary amount of energy to push it forward:

write 3 application of electrolysis

The same principle applies here: if we want NaCl to convert into Na and Cl 2 , we must provide the necessary energy, which in other words, would mean pushing the electrons in the opposite direction by taking them from Cl – ions and giving them to Na + ions.

It turns out that reversing the direction of a redox reaction is possible if external electrical power is used. In this case, we have an electrolytic cell where electrical current drives an otherwise nonspontaneous redox reaction through a process called electrolysis . So, let’s break down this process to gain a better understating using the sample of NaCl.

Electrolysis of Molten NaCl

The first thing we will need to do electrolysis of NaCl is a battery with an estimated potential of greater than 4.07 V as this was the estimated cell potential of the reaction between Na and Cl 2 . This cell potential is simply a qualitative description to emphasize that the reverse reaction will be nonspontaneous. Why qualitative? One reason would be the fact that we are going to need molten NaCl which is possible only at 800 o C and that is far from standard conditions. 

Molten NaCl is needed because if we use a solution of NaCl, there is also the possibility of the water undergoing electrolysis. 

We will discuss the electrolysis of aqueous NaCl as well, but for now, let’s draw what we have so for the electrolysis of molten NaCl:

write 3 application of electrolysis

The electrode on the left side is connected to the positive end of the battery and therefore, it is positively charged, and the one on the right side is negatively charged.

Notice that the electrodes themselves are not part of the reactions and therefore, unlike galvanic cells, the electrodes of an electrolytic cell are typically made of the same metal .

So, what is going to happen with the Na + and Cl – ions in the presence of these electrodes?

Because opposite charges attract, Na + ions are going to accumulate at the negatively charged electrode, and the Cl – ions will surround the positively charged electrode:

write 3 application of electrolysis

Now, the negatively charged electrode is going to push electrons to the Na + ions thus reducing them (cathode) to Na metal. On the other hand, the positively charged electrode needs electrons it is going to pull them from the Cl – ions despite the fact that they do not want to give them. This is the oxidation half-reaction and therefore, the electrode is the anode . The key here is the difference in the power of pulling the electrons and since the battery has a greater potential than the reaction between Na and Cl 2 , it does take the electrons from Cl – ions:

write 3 application of electrolysis

So, on the left electrode, Cl – ions are being oxidized to Cl 2 , and therefore, this is the anode , while the Na + ions are being reduced on the right side electrode making it the cathode .

Remember, regardless of the type of electrochemical cell, oxidation always occurs on the anode , and the reduction occurs on the cathode .

Notice, however, that the anode is positively charged here unlike what we saw in galvanic cells, where it was considered to be negatively charged as it gave the electrons to the cathode thus undergoing oxidation. In electrolytic cells, the external power source draws the electrons from the anode thus making it positive. So, keep in mind:

  • In a voltaic cell , the anode is negative as it takes the electrons from the cathode and thus making positive .
  • In an electrolytic cell , the anode is positive as the external battery takes the electrons from it, and provides them to the cathode making it negative .

Once again, in both cases, oxidation always takes place on the anode , and the reduction takes place on the cathode .

So, let’s write down each half-reaction. On the anode, we have the oxidation of Cl – to Cl 2 , and on the cathode, we have the reduction of Na + to Na:

Anode (oxidation): 2 Cl – ( l ) → Cl 2 ( g )  + 2 e –   

Cathode (reduction): 2Na + ( l ) + 2e- → 2Na( s )

The overall reaction then will be written as:

2Na + ( l ) + 2 Cl – ( l ) → 2Na( s ) + Cl 2 ( g ) 

Other Electrolysis Reactions

The electrolysis of NaCl is by far the most common example you will see in textbooks and during lectures, however, the principle of electrolysis can be applied to any nonspontaneous redox reaction . For example, Cu 2+ ions can oxidize Sn to Sn 2+ because of their higher reduction potential (greater tendency to be reduced) than Sn 2+ :

Sn( s ) + Cu 2+ ( aq ) → Sn 2+ ( aq ) + Cu( s )

The values of half-reactions can be found in a standard textbook, and to find the cell potential, we simply add them after switching the sign for the reverse reaction.

Sn( s ) → Sn 2+ ( aq ) + 2e –   E o  = +0.14 V

Cu 2+ ( aq ) + 2e –   → Cu( s )  E o  = +0.34 V ______________________________________

Sn( s ) + Cu 2+ ( aq ) → Sn 2+ ( aq ) + Cu( s )  E o  = +0.48 V

Feel free to check this article for more details on calculating the standard cell potential.

Now, if we set up an electrolytic cell, we can force the reverse reaction where Sn 2+ ions oxidize Cu to Cu 2+ ions granted the external voltage is higher than 0.48 V. With that said, it is not really Sn 2+  that oxidizes Cu, but rather the external voltage. The metals and their electrolyte simply support the current flow by providing and accepting the electrons. 

write 3 application of electrolysis

Electrolysis of Aqueous NaCl

Now that we have gone over the principle of electrolysis, let’s discuss perhaps a more confusing example which is the electrolysis of NaCl in an aqueous solution. The difference here compared to the electrolysis of molten sodium chloride is that the water can also be oxidized and/or reduced when external power is provided according to the following half-reactions:

Oxidation (anode): 2 H 2 O( l ) → O 2 ( g ) + 4 H + ( aq) + 4 e –    E ° = -1.23 V

Reduction (cathode): 2 H 2 O( l ) + 2 e – → H 2 ( g )  + 2 OH – ( aq )    E ° = -0.83 V

Note: The cell potentials given above are for standard conditions which imply 1 M H + . However, at pH 7, [H + ] = [OH – ] = 10 -7 M, so we do not have standard conditions, and the cell potentials are 0.82 V and -0.41 V respectively.

Now, each of these reactions is going to compete with the oxidation and reduction reactions that we saw in the electrolysis of molted NaCl. Let’s recall: Cl – ions were oxidized to Cl 2 according to the following equation:

2 Cl – ( aq ) → Cl 2 ( g )  + 2 e –

And Na + ions were reduced as they accept the electrons from the Cl – ions:

Na + ( aq ) + e- → Na( s )

So, the two possible oxidation reactions in the NaCl solution would be:

2 H 2 O( l ) → O 2 ( g ) + 4 H + ( aq ) + 4 e –    E ° = -1.23 V

2 Cl – ( aq ) → Cl 2 ( g )  + 2 e –    E ° = -1.36 V

Based on the higher cell potential, we’d expect water to be oxidized easier than Cl – . Interestingly though, it was found that Oxygen gas is not formed at the anode, and instead, chlorine gas is produced. How can we explain this? The answer is what’s called overvoltage or overpotential which is the difference between the electrode potential and the actual voltage required to cause electrolysis. The need for overpotential is caused by the slow electron transfer rate at the electrode–solution interface thus limiting the amount of current passing through an electrolytic cell. It turns out that the overpotential is higher for the formation of oxygen compared to that of chlorine and therefore, the oxidation of Cl – ions producing Cl 2 is what occurs at the anode.

Overpotential often makes it difficult to predict with certainty which of the possible half-reactions with close E ° values will occur, and only experiments can show what actually happens.

So, this was about the oxidation at the anode. The two possible reduction reactions at the cathode would be:

Na + ( aq ) + e- → Na( s )  E ° = -2.71 V

2 H 2 O( l )  + 2 e – → H 2 ( g )  + 2 OH – ( aq )  E ° = -0.83 V

The E o values are quite different, and as expected, because of a higher (less largely negative) standard reduction potential, water is reduced preferentially producing bubbles of hydrogen gas at the cathode.

Therefore, by adding the two half-reactions, we can obtain the overall electrolysis reaction:

Anode ( oxidation ) : 2 Cl – ( aq ) → Cl 2 ( g )  + 2 e –    E ° = -1.36 V

Cathode ( reduction ) : 2 H 2 O( l ) + 2 e – → H 2 ( g )  + 2 OH – ( aq )  E ° = -0.83 V

______________________________________________________________

2 Cl – ( aq ) + 2 H 2 O( l )  + 2 e – → Cl 2 ( g )  + 2 e –   + H 2 ( g )  + 2 OH – ( aq )

Overall cell reaction: 2 Cl – ( aq ) + 2 H 2 O( l ) → Cl 2 ( g ) + H 2 ( g )  + 2 OH – ( aq ) E ° = -2.19 V

This shows that, unlike in the electrolysis of molten NaCl, the Na + ions are simply spectator ions and do not participate in the electrolysis of aqueous NaCl.

To summarize, remember that electrolysis is performed in an electrolytic cell, and it is the process of driving a nonspontaneous redox reaction by using external electrical power.

  • Both in galvanic and electrolytic cells, oxidation always takes place on the anode , and the reduction takes place on the cathode .
  • In a voltaic cell , the anode is negative as it takes the electrons from the cathode and thus making it positive .
  • In an electrolytic cell , the anode is positive as the external battery takes the electrons from it and provides them to the cathode making it negative .

Electrolysis also has numerous other applications . For example, fluorine cannot be prepared by any spontaneous chemical reaction. It was found in 1886 by the French chemist Henri Moissan that passing electrical current through a molten mixture of potassium fluoride and hydrogen fluoride produces Fluorine, and it is still prepared commercially by the same method. Another application of electrolysis is the preparation of metals from their oxides which is how most metals are found in Earth’s crust.

In the next article, we will discuss the electrolysis of water.

  • Balancing Redox Reactions
  • Galvanic Cells
  • How to Calculate Standard Cell Potential
  • The Correlation Between  E o cell,  ΔG°, and  K
  • Nernst Equation
  • Nernst Equation Practice Problems
  • Concentration Cells
  • Electrolytic Cells
  • Electrolysis of Water
  • Calculating the Mass of Metal in Electroplating
  • Cell Potential Practice Problems
  • E o , ΔG o , K –  Practice Problems
  • Electrochemistry Practice Problems

Leave a Comment Cancel reply

Notify me of followup comments via e-mail. You can also subscribe without commenting.

web analytics

Logo for BCcampus Open Publishing

Want to create or adapt books like this? Learn more about how Pressbooks supports open publishing practices.

Chapter 14. Oxidation and Reduction

Electrolysis

Learning Objectives

  • Describe electrolysis from a perspective of redox reactions.
  • Give examples of electrolysis applications.

Up to this point, we have considered redox reactions for processes that are spontaneous. When set up as a voltaic cell or battery, such reactions can be used as a source of electricity. However, it is possible to go in the other direction. By forcing electricity into a cell, we can make a redox reaction occur that normally would not be spontaneous. Under these circumstances, the cell is called an electrolytic cell , and the process that occurs in the cell is called electrolysis  (Figure 14.4 “Electrolysis”).

Electrolysis

Electrolysis has many applications. For example, if NaCl is melted at about 800°C in an electrolytic cell and an electric current is passed through it, elemental sodium will appear at the cathode and elemental chlorine will appear at the anode as the following two reactions occur:

Na + + e − → Na 2 Cl − → Cl 2 + 2e −

Normally we expect elemental sodium and chlorine to react spontaneously to make NaCl. However, by using an input of electricity, we can force the opposite reaction to occur and generate the elements. Lithium, potassium, and magnesium can also be isolated from compounds by electrolysis.

Another element that is isolated by electrolysis is aluminum. Aluminum formerly was a difficult metal to isolate in its elemental form; in fact, the top of the Washington Monument has a 2.8 kg cap of aluminum metal, which at the time — 1884 — was the largest piece of elemental aluminum ever isolated. However, in 1886 the American Charles Hall and the Frenchman Paul Héroult almost simultaneously worked out an electrolytic process for isolating aluminum from bauxite, an ore of aluminum whose chemical formula is AlO x (OH) 3 − 2 x . The basic reactions are as follows:

Al 3+ + 3e − → Al 2O 2− → O 2 + 4e −

With the development of the Hall-Héroult process, the price of aluminum dropped by a factor of over 200, and aluminum metal became common. So much elemental aluminum is produced in the United States each year that it has been estimated that the electrolysis of aluminum uses 5% of all the electricity in the country. (Recycling aluminum requires about 1/70th the energy of refining aluminum from ore, which illustrates the tremendous energy savings that recycling provides.)

Another application of electrolysis is electroplating , which is the deposition of a thin layer of metal on an object for protective or decorative purposes (see Figure 14.5). Essentially, a metal object is connected to the cathode of an electrolytic cell and immersed in a solution of a particular metal cation. When the electrolytic cell is operated, a thin coating of the metal cation is reduced to the elemental metal on the surface of the object; the thickness of the coating can be as little as a few micrometers (10 −6 m). Jewellery, eating utensils, electrical contacts, and car parts like bumpers are common items that are electroplated. Gold, silver, nickel, copper, and chromium are common metals used in electroplating.

Test Cell

Key Takeaways

  • Electrolysis is the forcing of a nonspontaneous redox reaction to occur by the introduction of electricity into a cell from an outside source.
  • Electrolysis is used to isolate elements and electroplate objects.
  • Define electrolytic cell .
  • How does the operation of an electrolytic cell differ from a voltaic cell?
  • List at least three elements that are produced by electrolysis.
  • Write the half reactions for the electrolysis of the elements listed in Exercise 3.
  • Based on Table 14.1 “Standard Reduction Potentials of Half Reactions” , what voltage must be applied to an electrolytic cell to electroplate copper from Cu 2+ ?
  • Based on Table 14.1, what voltage must be applied to an electrolytic cell to electroplate aluminum from Al 3+ ?
  • An electrochemical cell in which charge is forced through and a nonspontaneous reaction occurs.
  • Any three of the following: Al, K, Li, Na, Cl2, or Mg

Media Attributions

  • “Hullcell” © 2009 by Fstep is licensed under a Public Domain license

Introductory Chemistry - 1st Canadian Edition Copyright © 2014 by Jessie A. Key is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License , except where otherwise noted.

Share This Book

write 3 application of electrolysis

Chapter 14. Oxidation and Reduction

Electrolysis, learning objectives.

  • Describe electrolysis from a perspective of redox reactions.
  • Give examples of electrolysis applications.

Up to this point, we have considered redox reactions for processes that are spontaneous. When set up as a voltaic cell or battery, such reactions can be used as a source of electricity. However, it is possible to go in the other direction. By forcing electricity into a cell, we can make a redox reaction occur that normally would not be spontaneous. Under these circumstances, the cell is called an electrolytic cell , and the process that occurs in the cell is called electrolysis  ( Figure 14.5 “Electrolysis” ).

Figure 14.5 Electrolysis

Electrolysis

In an electrolytic cell, electricity is forced through the cell to induce a nonspontaneous redox reaction. Here, the redox reaction 2 H 2 O → 2 H 2 + O 2 is being caused by the introduction of electricity, which is supplied by the battery.

Electrolysis has many applications. For example, if NaCl is melted at about 800°C in an electrolytic cell and an electric current is passed through it, elemental sodium will appear at the cathode and elemental chlorine will appear at the anode as the following two reactions occur:

Na + + e − → Na 2 Cl − → Cl 2 + 2e −

Normally we expect elemental sodium and chlorine to react spontaneously to make NaCl. However, by using an input of electricity, we can force the opposite reaction to occur and generate the elements. Lithium, potassium, and magnesium can also be isolated from compounds by electrolysis.

Another element that is isolated by electrolysis is aluminum. Aluminum formerly was a difficult metal to isolate in its elemental form; in fact, the top of the Washington Monument has a 2.8 kg cap of aluminum metal, which at the time—1884—was the largest piece of elemental aluminum ever isolated. However, in 1886 the American Charles Hall and the Frenchman Paul Héroult almost simultaneously worked out an electrolytic process for isolating aluminum from bauxite, an ore of aluminum whose chemical formula is AlO x (OH) 3 − 2 x . The basic reactions are as follows:

Al 3+ + 3e − → Al 2 O 2− → O 2 + 4e −

With the development of the Hall-Héroult process, the price of aluminum dropped by a factor of over 200, and aluminum metal became common. So much elemental aluminum is produced in the United States each year that it has been estimated that the electrolysis of aluminum uses 5% of all the electricity in the country. (Recycling aluminum requires about 1/70th the energy of refining aluminum from ore, which illustrates the tremendous energy savings that recycling provides.)

Another application of electrolysis is electroplating , which is the deposition of a thin layer of metal on an object for protective or decorative purposes ( Figure 14.6 ). Essentially, a metal object is connected to the cathode of an electrolytic cell and immersed in a solution of a particular metal cation. When the electrolytic cell is operated, a thin coating of the metal cation is reduced to the elemental metal on the surface of the object; the thickness of the coating can be as little as a few micrometers (10 −6 m). Jewelry, eating utensils, electrical contacts, and car parts like bumpers are common items that are electroplated. Gold, silver, nickel, copper, and chromium are common metals used in electroplating.

Figure 14.6

Test Cell

Source: Photo courtesy of Fstep, http://commons.wikimedia.org/wiki/File:Hullcell.jpg .

Key Takeaways

  • Electrolysis is the forcing of a nonspontaneous redox reaction to occur by the introduction of electricity into a cell from an outside source.
  • Electrolysis is used to isolate elements and electroplate objects.

Define electrolytic cell .

How does the operation of an electrolytic cell differ from a voltaic cell?

List at least three elements that are produced by electrolysis.

Write the half reactions for the electrolysis of the elements listed in Exercise 3.

Based on Table 14.1 “Standard Reduction Potentials of Half Reactions” , what voltage must be applied to an electrolytic cell to electroplate copper from Cu 2+ ?

Based on Table 14.1 “Standard Reduction Potentials of Half Reactions” , what voltage must be applied to an electrolytic cell to electroplate aluminum from Al 3+ ?

an electrochemical cell in which charge is forced through and a nonspontaneous reaction occurs

any three of the following: Al, K, Li, Na, Cl 2 , or Mg

  • Introductory Chemistry- 1st Canadian Edition . Authored by : Jessie A. Key and David W. Ball. Provided by : BCCampus. Located at : https://opentextbc.ca/introductorychemistry/ . License : CC BY-NC-SA: Attribution-NonCommercial-ShareAlike . License Terms : Download this book for free at http://open.bccampus.ca

Footer Logo Lumen Candela

Privacy Policy

Get verified solutions to all your doubts

35 lakh questions, verified answers.

Industrial Applications of Electrolysis

Electrolysis is used in industry for a wide range of applications. Although the individual setups of the electrolytic cell may differ for each of these applications, the core process is still the same.

Extracting and refining of metals

Electrolysis can be used to extract metal from its ore . This process is used for metal ores containing metals such as aluminium and sodium. Also, electrolysis can be used to refine metals such as copper.

Production of chemicals

The chemical reactions in electrolysis can produce useful chemicals. For example, electrolysis of sodium chloride solution, also known as brine, can produce sodium hydroxide, hydrogen, and chlorine. These chemicals have a variety of uses in manufacturing and other applications. Another example of this application is the electrolysis of water to produce hydrogen and oxygen.

Electroplating

Electroplating is the process of coating an object in metal. This process can be used to give an object a decorative finish, protect the object from corrosion, or give an object new material properties. Electroplating is used in the creation of jewelry, kitchen utensils, automotive parts, and components in electrical circuits.

Talk to our experts

1800-120-456-456

Applications of Electrolysis for IIT JEE

  • Applications Of Electrolysis

ffImage

What is Electrolysis?

Electrolysis is a process of chemical decomposition produced by passing an electric current through a liquid or solution containing ions. In this process, we generally pass a direct electrical current through an ionic substance which is either present in its molten form or is dissolved in any appropriate solvent, producing chemical reactions at the electrodes and causing material decomposition. The key process of electrolysis is the exchange of atoms and ions from the external circuit by removing or adding electrons. Often the desired electrolysis products are in a different physical state from the electrolyte and some physical processes can remove them. The main components needed for electrolysis are:

An Electrolyte: A substance, often an ion-conducting polymer that contains free ions in the electrolyte that carries electrical current. If the ions, as in most solid salts, are not mobile, electrolysis will not be possible. 

Direct Current (DC) Electrical Supply: It provides the necessary energy to generate or discharge the electrolyte ions. Electrical current is transported in the external circuit by electrons.

Two Electrodes: These are electrical conductors providing the physical interface between the electrolyte and the energy-providing electrical circuit.

1. Metal Extraction (Electrometallurgy)

Electrometallurgy is the process of electrolysis for extracting metal from the ore. Electrolytes of fused electrolytes obtain metals such as sodium, potassium, magnesium, calcium aluminum, etc. There are two methods of metal extraction based on ore’s physical state. The ore is treated with a strong acid in the first method to obtain a salt and such salt solution is electrolyzed to release the metal. The ore is in a molten state in the second process and electrolysed in a furnace.

(i) Extraction of Zinc:

Sulphuric acid is used to treat the zinc ore. During the electrolysis process involved in this reaction, the zinc sulphate solution is obtained as a result, which also acts as an electrolyte for the electrolytic reaction. The cathodes are made of aluminum in the electrolyte tank and the anodes are made of lead. The current density is maintained at 1000 A / m² and the cathode deposits zinc. The consumption of energy is between 3000 and 5000 KWH per tonne.

(ii) Extraction of Aluminium:

Bauxite and cryolite are aluminum ores. These are chemically treated and reduced to aluminum oxide, then dissolved in fused cryolite and electrolysed in a furnace. Deposits of aluminum settle down at the bottom. The furnace temperature is about 1000 °C to keep the electrolyte in a fused state. To complete the process, the current of around 4000 ampere is needed and the energy consumed is around 20,000 to 25,000 kWH per tonne.

2. Production of Non - Metals

Electrolysis is used to obtain non-metals such as hydrogen, fluorine, chlorine. Hydrogen is obtained in the presence of electrolytes such as H₂SO₄ and KNO₃ through the electrolysis of water.

3. Metal Refining

The main advantages of electrolytic processes of extracting a metal are that the purity of the product obtained is between 98 and 99 %. Electrolysis is used for further refining. The anode is made of metal that has been extracted. The cathode deposits pure metal. The electrolyte is made of metal solution such as copper, copper sulphate and nickel, nickel chloride. Electrolytic process energy consumption in copper refining is 150 to 300 kWH / ton of refined copper. Copper purity can be enhanced by electro-refining. The same electroplating or electrodeposition principle is used in this case. The electrolyte is a copper ion solution, like copper sulphate. The cathode is a strip of pure copper and the anode is an unclean copper lump. At the anode: copper atoms dissolve to form copper ions from the impure copper. This causes the size of the anode to decline. Impurities are gathered in the cell and fall off the anode. At the cathode: ions Cu2 + migrate to the cathode where they are deposited as solid copper on the cathode. The cathode is getting thicker (increasing in size).

4. Production of Chemicals

Many chemicals like caustic soda (NaOH) chlorine gas etc. are produced on a large scale through electrolysis. Large-scale electrolysis also produces potassium permanganate, hydrogen and oxygen etc.

5. Electroplating

Electroplating is an object's coating with a thin metal film deposited during electrolysis by an electrolytic solution. Electroplating is often used on items such as knives and forks (silver plate) to produce an attractive and durable finish. It is also used to protect metals susceptible to corrosion. For example, relatively unreactive cadmium metal is often placed on steel components to inhibit rusting. Electroplating is a very valuable industrial process, but its use requires expensive and consistently efficient treatment of the waste that it produces. It is possible to produce inexpensive silver-plated jewelry by electroplating. Gold rings, which cause the fingers to turn green, are in fact gold-electroplated copper rings.

The article to be plated is used as the cathode in the electroplating process and the metal plated on the article is used as the anode. The electrolytic solution or ‘bath’ contains a plated metal salt. A low-voltage electrical current causes metal ions from the bath to gain electrons from the cathode and deposit on the cathode (the object) as a metal coating. It also causes metal atoms to lose electrons on the anode and go as ions into the bath. The anode gradually disappears as the plating proceeds and maintains the concentration of the metal ion in the bath. Many toxic solutions are used in electroplating operations. Objects to be plated with concentrated acidic or basic solutions are thoroughly cleaned of all grease and dirt. Eventually, due to contamination, cleaning solutions become ineffective and must be disposed of.

A number of factors contribute to the quality of the metal coating formed by electroplating:

The concentration of the cations to be reduced is carefully controlled in the plating solution. It is necessary to avoid unwanted side reactions.

The electrolyte type and concentration should be carefully considered.

Compounds to control acidity and increase conductivity must be included in the solution.

Some compounds make it brighter or smoother to cover the metal.

The anode must be shaped like a cathode object in many electroplating cells to achieve even metal coating.

6. Electrotyping

This is a process by which a process of electro-plating reproduces wood cuts etc. in copper. A mould is first made of the type in wax in this process, then coated with black lead to give it a metallic surface and then subjected to the electro-deposition process. Thus, on the prepared surface, a copper film is formed.

7. Electroforming

This is another electro-deposition application. Electro-deposition reproduction of objects on some kind of mould or shape is known as electro-forming. This is another useful example among many electrolysis applications. We must take the impression of objects on wax or other wax-like material first. To make it conductive, the surface of the wax mold that bears the object's exact impression is coated with graphite powder. Then the mold is immersed as a cathode in the electrolyte solution. The electrolyte metal will be deposited on the impressed graphite surface of the mold during the electrolysis process. The article will be removed after obtaining a layer of desired thickness and the wax will be melted to obtain the reproduced object as a metal shell. Reproduction of gramophone record disc is a popular use of electroforming. The original recording takes place on a wax composition record. This wax mold is then coated to make it conductive with gold powder. This mold is then immersed as a cathode in a blue vitriol electrolyte. Using a copper anode keeps the solution saturated. On the wax mold, the copper electroforming produces a master plate used to stamp a large number of shellac disks.

8. Electrocleaning

The item to be cleaned from oil and grease is the cathode and the solution is passed through the iron tank or vat filled with an electrolyte solution and heavy current. At cathode, caustic soda and hydrogen are produced which removes the grease from the article surface. The process is known as cathodic cleaning and is applied to zinc and aluminum. Article is made anode for anodic cleaning.

9. Determination of Equivalent Elemental Masses

The quantities of metals deposited on the cathodes of the two cells are proportional to their equivalent masses of the respective metals, according to the second law of electrolysis when the same quantity of electronic current is passed through salt solutions of two different cells. If the quantities of metals deposited on the cathodes are WA and WB, then 

\[\frac{W_{A}}{W_{B}} = \frac{Equivalent\:Mass\: of\: A}{Equivalent\: Mass\: of\: B}\]

Knowing the equivalent mass of one metal, it is possible to calculate the equivalent mass of the other metal from the relationship above. This method can also determine the equivalent masses of those non-metals that evolve at anodes.

10. Thickness of any Coated Layer

Electrolysis method can also be used to calculate the thickness of the coated layer of any material. For that we just need to know the dimensions of the coated material and we will get its thickness.

Metal Cleaning - Dirty metals can be cleaned by electrolysis process. The unclean metal is made of anode. The same metal plate is purely made of cathode. The salt solution is used as an electrolyte. When a current passes through an electrolyte, the anode usually melts. Pure metal is placed in the cathode and dirt is left behind.

Chemical Production- Caustic soda is prepared by electrolysis of sodium chloride solution.

Anodising- It is the process of coating aluminium with its oxide to protect it from corrosion.

Production of O 2 and H 2 - These are obtained by electrolysis of acidic water.

Medical Applications- Electrolysis is used to rejuvenate polio; removing unwanted hair from any part of the body.

Electrolytic Process

In the process of electrolysis, ions exchange with atoms due to the addition or removal of electrons from the outer circle. Basically, in the current passage, the cations go to the cathode, take the electrons from the cathode (given the source-battery source), and release them from a neutral atom. A neutral atom, if solid, is placed in the cathode and, if gas, rises to the surface. This is a reduction process and the cation decreases in the cathode. At the same time, the anions, in turn, supply their electrons to the anode and are oxidized by neutral atoms in the anode. The electrons released by the anions travel across the electrical circuit and reach the cathode completing the circuit. Electrolysis combines a simultaneous oxidation reaction at an anode with a reduction reaction at the cathode.

For example, when an electric current, passing through a molten sodium chloride, the sodium ion is attracted to the cathode, from which it emits an electrode and becomes a sodium atom. The Electrolysis process, although useful for obtaining basic forms from computers directly, can also be applied indirectly to alkali and alkaline earth metals, metal cleaning, metal insertion, compound preparation etc.

arrow-right

FAQs on Applications of Electrolysis for IIT JEE

1. What is the basic meaning of electrolysis?

Electrolysis is that branch of chemistry that is responsible for the chemical reactions that occur in electrical energy and the production of electricity through chemical reactions. An electrochemical cell consists of 2 metal conductors known as electrodes connected to an ionic conductor i.e. electrolyte. The electrolyte and electrode form the Electrode Compartment. Electrolytic cells are those where the automatic reaction is controlled by an external current source of electricity. Galvanic cells produce electrical energy due to the automatic reaction of cells.

2. What are the key topics in electrolysis?

Electrolysis consists of the underlying key topics, which will help you to obtain the most out of JEE. In electrolysis, there is a direct relationship between the amount of electricity flowing through a cell and the number of chemical reactions that occur. Here are some ideas to read carefully:

Electrolysis and electrolytic cell

Daniel cell

Electrochemical series

Electrode Power

Arrhenius Theory for Electrolytic Dissociation

Faraday's Law of Electrolysis

Nernst equation

Electrolytic Processing

Kohlrausch's Law

Focus cells

Use of Electrolysis

3. Is electrolysis used in metallurgical processes?

Yes, electrolysis is widely utilized in metallurgical processes in a good amount. Electrolysis is widely used in metallurgical processes, such as the extraction (electromining) or refining (electrorefining) of ore or compounds and in the incorporation of metals from the solution (electroplating). Metallic sodium and chlorine gas are produced by the electrolysis of molten sodium chloride; Electrolysis of aqueous sodium chloride solution produces sodium hydroxide and chlorine gas. Hydrogen and oxygen are produced by water electrolysis. Products can be neutral elements or new molecules.

4. How can you download the applications of electrolysis in IIT JEE from online?

To download the study material resources you just have to search for what you require. After finding the appropriate and legible study material you have to click on the link and download the study material. Vedantu is one of the best websites to download the study material for the Applications of Electrolysis in IIT JEE. The students of IIT JEE must be using the link to the website of Vedantu from which they will get the study material. Guaranteed by experts and tested by many other students Vedantu provides the best study material.

5. Is vedantu a reliable website for referring to the applications of electrolysis in IIT JEE?

You should learn from Vedantu as it is the only educational website that expands student horizons by providing realistic and focused solutions to building students ’minds and creating better understanding.  A student must look into the topic of Applications of Electrolysis prepared according to the understanding level of the students from the experts of Vedantu. Vedantu closes the gap between students and the most important aspect, complete learning. Finally, Vedantu solutions are fully available for free on the Vedantu website. Students will not face a problem.

Logo for University of Central Florida Pressbooks

Chapter 17: Electrochemistry

17.6 Electrolysis

Learning outcomes.

  • Describe the process of electrolysis
  • Compare the operation of electrolytic cells with that of galvanic cells
  • Perform stoichiometric calculations for electrolytic processes

Electrochemical cells in which spontaneous redox reactions take place ( galvanic cells ) have been the topic of discussion so far in this chapter. In these cells, electrical work is done by a redox system on its surroundings as electrons produced by the redox reaction are transferred through an external circuit. This final section of the chapter will address an alternative scenario in which an external circuit does work on a redox system by imposing a voltage sufficient to drive an otherwise nonspontaneous reaction, a process known as electrolysis . A familiar example of electrolysis is recharging a battery, which involves use of an external power source to drive the spontaneous (discharge) cell reaction in the reverse direction, restoring to some extent the composition of the half-cells and the voltage of the battery. Perhaps less familiar is the use of electrolysis in the refinement of metallic ores, the manufacture of commodity chemicals, and the electroplating of metallic coatings on various products (e.g., jewelry, utensils, auto parts). To illustrate the essential concepts of electrolysis, a few specific processes will be considered.

The Electrolysis of Molten Sodium Chloride

Metallic sodium, Na, and chlorine gas, [latex]\ce{Cl2}[/latex], are used in numerous applications, and their industrial production relies on the large-scale electrolysis of molten sodium chloride, NaCl( l ). The industrial process typically uses a Downs cell similar to the simplified illustration shown in Figure 17.6.1. The reactions associated with this process are:

[latex]\text{anode:} \,\,\,\, \ce{2Cl-} (l) \longrightarrow \ce{Cl2} (g) + \ce{2e-}[/latex][latex]\text{cathode:} \,\,\,\, \ce{Na+} (l) + \ce{e-} \longrightarrow \ce{Na} (l)[/latex] [latex]\text{cell:} \,\,\,\, \ce{2Na+} (l) + \ce{2Cl-} (l) \longrightarrow \ce{2Na} (l) + \ce{Cl2} (g)[/latex]

The cell potential for the above process is negative, indicating the reaction as written (decomposition of liquid [latex]\ce{NaCl}[/latex]) is not spontaneous. To force this reaction, a positive potential of magnitude greater than the negative cell potential must be applied to the cell.

This diagram shows a tank containing a light blue liquid, labeled “Molten N a C l.” A vertical dark grey divider with small, evenly distributed dark dots, labeled “Porous screen” is located at the center of the tank dividing it into two halves. Dark grey bars are positioned at the center of each of the halves of the tank. The bar on the left, which is labeled “Anode” has green bubbles originating from it. The bar on the right which is labeled “Cathode” has light grey bubbles originating from it. An arrow points left from the center of the tank toward the anode, which is labeled “C l superscript negative.” An arrow points right from the center of the tank toward the cathode, which is labeled “N a superscript plus.” A line extends from the tops of the anode and cathode to a rectangle centrally placed above the tank which is labeled “Voltage source.” An arrow extends upward above the anode to the left of the line which is labeled “e superscript negative.” A plus symbol is located to the left of the voltage source and a negative sign it located to its right. An arrow points downward along the line segment leading to the cathode. This arrow is labeled “e superscript negative.” The left side of below the diagram is the label “2 C l superscript negative right pointing arrow C l subscript 2 ( g ) plus 2 e superscript negative.” At the right, below the diagram is the label “2 N a superscript plus plus 2 e superscript negative right pointing arrow 2 N a ( l ).”

The Electrolysis of Water

Water may be electrolytically decomposed in a cell similar to the one illustrated in Figure 17.6.2 . To improve electrical conductivity without introducing a different redox species, the hydrogen ion concentration of the water is typically increased by addition of a strong acid. The redox processes associated with this cell are

[latex]\text{anode:} \,\,\,\, \ce{2H2O} (l) \longrightarrow \ce{O2} (g)+ \ce{4H+} (aq) + \ce{4e-} \,\,\,\,\, \text{E}^{\circ}_\text{anode} = +1.229\text{V}[/latex][latex]\text{cathode:} \,\,\,\, \ce{2H+} (aq) + \ce{2e-} \longrightarrow \text{H2}(g) \,\,\,\,\, \text{E}^{\circ}_\text{cathode} = 0\text{V}[/latex] [latex]\text{cell:} \,\,\,\, \ce{2H2O} (l) \longrightarrow \ce{2H2} (g) + \ce{O2} (g) \,\,\,\,\, \text{E}^{\circ}_\text{cell} = -1.229\text{V}[/latex]

Again, the cell potential as written is negative, indicating a nonspontaneous cell reaction that must be driven by imposing a cell voltage greater than +1.229 V. Keep in mind that standard electrode potentials are used to inform thermodynamic predictions here, though the cell is not operating under standard state conditions. Therefore, at best, calculated cell potentials should be considered ballpark estimates.

This figure shows an apparatus used for electrolysis. A central chamber with an open top has a vertical column extending below that is nearly full of a clear, colorless liquid, which is labeled “H subscript 2 O plus H subscript 2 S O subscript 4.” A horizontal tube in the apparatus connects the central region to vertical columns to the left and right, each of which has a valve or stopcock at the top and a stoppered bottom. On the left, the stopper at the bottom has a small brown square connected just above it in the liquid. The square is labeled “Anode plus.” A black wire extends from the stopper at the left to a rectangle which is labeled “Voltage source” on to the stopper at the right. The left side of the rectangle is labeled with a plus symbol and the right side is labeled with a negative sign. The stopper on the right also has a brown square connected to it which is in the liquid in the apparatus. This square is labeled “Cathode negative.” The level of the solution on the left arm or tube of the apparatus is significantly higher than the level of the right arm. Bubbles are present near the surface of the liquid on each side of the apparatus, with the bubbles labeled as “O subscript 2 ( g )” on the left and “H subscript 2 ( g )” on the right.

The Electrolysis of Aqueous Sodium Chloride

When aqueous solutions of ionic compounds are electrolyzed, the anode and cathode half-reactions may involve the electrolysis of either water species ([latex]\ce{H2O}[/latex], [latex]\ce{H+}[/latex], [latex]\ce{OH-}[/latex]) or solute species (the cations and anions of the compound). As an example, the electrolysis of aqueous sodium chloride could involve either of these two anode reactions:

(i) [latex]\ce{2Cl-}(aq) \longrightarrow \ce{Cl2}(g) + \ce{2e-} \,\,\,\,\,\,\, \text{E}^\circ_{\text{anode}} = +1.35827 \text{V}[/latex]

(ii) [latex]\ce{2H2O}(l) \longrightarrow \ce{O2}(g) + \ce{4H+}(aq) + \ce{4e-}\,\,\,\,\,\,\,\text{E}^\circ_{\text{anode}}= +1.229 \text{V}[/latex]

The standard electrode ( reduction ) potentials of these two half-reactions indicate water may be oxidized at a less negative/more positive potential (–1.229 V) than chloride ion (–1.358 V). Thermodynamics thus predicts that water would be more readily oxidized, though in practice it is observed that both water and chloride ion are oxidized under typical conditions, producing a mixture of oxygen and chlorine gas.

Turning attention to the cathode, the possibilities for reduction are:

(iii) [latex]\ce{2H+}(aq) + \ce{2e-} \longrightarrow \ce{H2}(g) \,\,\,\,\,\,\, \text{E}^\circ_{\text{cathode}}= 0 \text{V}[/latex]

(iv) [latex]\ce{2H2O}(l) + \ce{2e-} \longrightarrow \ce{H2}(g) + \ce{2OH-}(aq) \,\,\,\,\,\,\, \text{E}^\circ_{\text{cathode}}= −0.8277 \text{V}[/latex]

(v) [latex]\ce{Na+}(aq) + \ce{e-} \longrightarrow \ce{Na}(s) \,\,\,\,\,\,\, \text{E}^\circ_{\text{cathode}}= −2.71 \text{V}[/latex]

Comparison of these standard half-reaction potentials suggests the reduction of hydrogen ion is thermodynamically favored. However, in a neutral aqueous sodium chloride solution, the concentration of hydrogen ion is far below the standard state value of 1 M (approximately 10 -7 M ), and so the observed cathode reaction is actually reduction of water. The net cell reaction in this case is then

[latex]\text{cell: } \ce{2H2O}(l) + \ce{Cl-}(aq) \longrightarrow \ce{H2}(g) + \ce{Cl2}(g) + \ce{2OH-}(aq) \,\,\,\,\,\,\, \text{E}^\circ_{\text{cell}} = −2.186 \text{V}[/latex]

This electrolysis reaction is part of the chlor-alkali process used by industry to produce chlorine and sodium hydroxide (lye).

Electroplating

An important use for electrolytic cells is in electroplating . Electroplating results in a thin coating of one metal on top of a conducting surface. Reasons for electroplating include making the object more corrosion resistant, strengthening the surface, producing a more attractive finish, or for purifying metal. The metals commonly used in electroplating include cadmium, chromium, copper, gold, nickel, silver, and tin. Common consumer products include silver-plated or gold-plated tableware, chrome-plated automobile parts, and jewelry. We can get an idea of how this works by investigating how silver-plated tableware is produced ( Figure 17.6.3 ).

This figure contains a diagram of an electrochemical cell. One beakers is shown that is just over half full. The beaker contains a clear, colorless solution that is labeled “A g N O subscript 3 ( a q ).” A silver strip is mostly submerged in the liquid on the left. This strip is labeled “Silver (anode).” The top of the strip is labeled with a red plus symbol. An arrow points right from the surface of the metal strip into the solution to the label “A g superscript plus” to the right. A spoon is similarly suspended in the solution and is labeled “Spoon (cathode).” It is labeled with a black negative sign on the tip of the spoon’s handle above the surface of the liquid. An arrow extends from the label “A g superscript plus” to the spoon on the right. A wire extends from the top of the spoon and the strip to a rectangle labeled “Voltage source.” An arrow points upward from silver strip which is labeled “e superscript negative.” Similarly, an arrow points down at the right to the surface of the spoon which is also labeled “e superscript negative.” A plus sign is shown just outside the voltage source to the left and a negative is shown to its right.

In Figure 17.6.3 , the anode consists of a silver electrode, shown on the left. The cathode is located on the right and is the spoon, which is made from inexpensive metal. Both electrodes are immersed in a solution of silver nitrate. Applying a sufficient potential results in the oxidation of the silver anode

[latex]\text{anode: } \ce{Ag}(s) \longrightarrow \ce{Ag+}(aq) + \ce{e-}[/latex]

and reduction of silver ion at the (spoon) cathode:

[latex]\text{cathode: } \ce{Ag+}(aq) + \ce{e-} \longrightarrow \text{Ag}(s)[/latex]

The net result is the transfer of silver metal from the anode to the cathode. Several experimental factors must be carefully controlled to obtain high-quality silver coatings, including the exact composition of the electrolyte solution, the cell voltage applied, and the rate of the electrolysis reaction (electrical current).

Quantitative Aspects of Electrolysis

Electrical current is defined as the rate of flow for any charged species. Most relevant to this discussion is the flow of electrons. Current is measured in a composite unit called an ampere, defined as one coulomb per second (A = 1 C/s). The charge transferred, Q , by passage of a constant current, [latex]I[/latex], over a specified time interval, [latex]t[/latex], is then given by the simple mathematical product

[latex]Q =It[/latex]

When electrons are transferred during a redox process, the stoichiometry of the reaction may be used to derive the total amount of (electronic) charge involved. For example, the generic reduction process

[latex]\text{M}^{\text{n}+} (aq) + \text{ne}^- \longrightarrow \text{M} (s)[/latex]

involves the transfer of n mole of electrons. The charge transferred is, therefore,

[latex]Q =nF[/latex]

where F is Faraday’s constant, the charge in coulombs for one mole of electrons. If the reaction takes place in an electrochemical cell, the current flow is conveniently measured, and it may be used to assist in stoichiometric calculations related to the cell reaction.

Example 17.6.1: Converting Current to Moles of Electrons

In one process used for electroplating silver, a current of 10.23 A was passed through an electrolytic cell for exactly 1 hour. How many moles of electrons passed through the cell? What mass of silver was deposited at the cathode from the silver nitrate solution?

Faraday’s constant can be used to convert the charge ( Q ) into moles of electrons ( n ). The charge is the current ( I ) multiplied by the time

[latex]\displaystyle{n}=\frac{Q}{F}=\frac{\frac{10.23\text{ C}}{\text{s}}\times\text{1 hr}\times\frac{60\text{ min}}{\text{hr}}\times\frac{60\text{ s}}{\text{min}}}{\text{96,485 C/mol e}^-}=\frac{\text{36,830 C}}{\text{96,485 C/mol e}^-}=0.3817\text{ mol e}^-[/latex]

From the problem, the solution contains [latex]\ce{AgNO3}[/latex], so the reaction at the cathode involves 1 mole of electrons for each mole of silver

[latex]\text{cathode:} \ce{Ag}^+(aq) + \ce{e-} \longrightarrow \ce{Ag}(s)[/latex]

The atomic mass of silver is 107.9 g/mol, so

[latex]\text{mass } \ce{Ag} = 0.3817\text{ mol e}^-\times\dfrac{1\text{ mol } \ce{Ag}}{1\text{ mol e}^-}\times\dfrac{107.9\ce{ g } \ce{Ag}}{1\text{ mol Ag}}=41.19\text{ g } \ce{Ag}[/latex]

Check your answer: From the stoichiometry, 1 mole of electrons would produce 1 mole of silver. Less than one-half a mole of electrons was involved and less than one-half a mole of silver was produced.

Check Your Learning

Example 17.6.2: time required for deposition.

In one application, a 0.010-mm layer of chromium must be deposited on a part with a total surface area of 3.3 m 2 from a solution of containing chromium(III) ions. How long would it take to deposit the layer of chromium if the current was 33.46 A? The density of chromium (metal) is 7.19 g/cm 3 .

This problem brings in a number of topics covered earlier. An outline of what needs to be done is:

  • If the total charge can be determined, the time required is just the charge divided by the current
  • The total charge can be obtained from the amount of [latex]\ce{Cr}[/latex] needed and the stoichiometry
  • The amount of [latex]\ce{Cr}[/latex] can be obtained using the density and the volume Cr required
  • The volume [latex]\ce{Cr}[/latex] required is the thickness times the area

Solving in steps, and taking care with the units, the volume of [latex]\ce{Cr}[/latex] required is

[latex]\text{volume}=\left(\text{0.010 mm}\times\frac{\text{1 cm}}{\text{10 mm}}\right)\times\left(\text{3.3 m}^{2}\times\left(\frac{\text{10,000 cm}^{2}}{\text{1 m}^2}\right)\right)=\text{33 cm}^3[/latex]

Cubic centimeters were used because they match the volume unit used for the density. The amount of [latex]\ce{Cr}[/latex] is then

mass = volume × density = [latex]33\cancel{\text{ cm}^3}\times\dfrac{7.19\text{ g}}{\cancel{\text{cm}^3}}=237\text{ g  } \ce{Cr}[/latex] mol Cr [latex]= 237\text{ g } \ce{Cr}\times\dfrac{1\text{ mol } \ce{Cr}}{52.00\text{g } \ce{Cr}}=4.56\text{ mol  } \ce{Cr}[/latex]

Since the solution contains chromium(III) ions, 3 moles of electrons are required per mole of [latex]\ce{Cr}[/latex]. The total charge is then

[latex]Q=4.56\text{ mol } \ce{Cr}\times\dfrac{3\text{ mol e}^-}{1\text{ mol } \ce{Cr}}\times\dfrac{96485\ce{ C}}{\text{mol e}^-}=1.32\times{10^{6}}\ce{C}[/latex]

The time required is then

[latex]t=\frac{Q}{I}=\dfrac{1.32\times10^6\text{ C}}{33.46\text{ C/s}}=3.95\times10^4\text{ s}=11.0\text{hr}[/latex]

Check your answer: In a long problem like this, a single check is probably not enough. Each of the steps gives a reasonable number, so things are probably correct. Pay careful attention to unit conversions and the stoichiometry.

Key Concepts and Summary

Nonspontaneous redox processes may be forced to occur in electrochemical cells by the application of an appropriate potential using an external power source—a process known as electrolysis. Electrolysis is the basis for certain ore refining processes, the industrial production of many chemical commodities, and the electroplating of metal coatings on various products. Measurement of the current flow during electrolysis permits stoichiometric calculations.

Key Equations

  • [latex]Q = I \times{t} = n\times{F}[/latex]
  • [latex]\ce{CaCl2}[/latex]
  • [latex]\ce{LiH}[/latex]
  • [latex]\ce{AlCl3}[/latex]
  • [latex]\ce{CrBr3}[/latex]
  • What mass of each product is produced in each of the electrolytic cells of the previous problem if a total charge of 3.33 × 105 [latex]\ce{C}[/latex] passes through each cell? Assume the voltage is sufficient to perform the reduction.
  • [latex]\ce{Al^3+}[/latex], 1.234 A
  • [latex]\ce{Ca^2+}[/latex], 22.2 A
  • [latex]\ce{Cr^5+}[/latex], 37.45 A
  • [latex]\ce{Au^3+}[/latex], 3.57 A
  • A current of 2.345 A passes through the cell shown in Figure for 45 minutes. What is the volume of the hydrogen collected at room temperature if the pressure is exactly 1 atm? Assume the voltage is sufficient to perform the reduction. (Hint: Is hydrogen the only gas present above the water?)
  • An irregularly shaped metal part made from a particular alloy was galvanized with zinc using a [latex]\ce{Zn(NO3)2}[/latex] solution. When a current of 2.599 A was used, it took exactly 1 hour to deposit a 0.01123-mm layer of zinc on the part. What was the total surface area of the part? The density of zinc is 7.140 g/cm 3 . Assume the efficiency is 100%.
  • mass [latex]\ce{Ca}[/latex] = 69.1 g ; mass [latex]\ce{Cl2}[/latex]   = 122 g
  • mass [latex]\ce{Li}[/latex] = 23.9 g ; mass [latex]\ce{H2}[/latex] = 3.48 g
  • mass [latex]\ce{Al}[/latex] = 31.0 g ; mass [latex]\ce{Cl2}[/latex] = 122 g
  • mass [latex]\ce{Cr}[/latex] = 59.8 g ; mass [latex]\ce{Br2}[/latex]   =276 g

electrolysis: process using electrical energy to cause a nonspontaneous process to occur

electrolytic cell: electrochemical cell in which electrolysis is used; electrochemical cell with negative cell potentials

electroplating: depositing a thin layer of one metal on top of a conducting surface

overpotential: difference between the theoretical potential and actual potential in an electrolytic cell; the “extra” voltage required to make some nonspontaneous electrochemical reaction to occur

CC licensed content, Shared previously

  • Chemistry 2e. Provided by : OpenStax. Located at : https://openstax.org/ . License : CC BY: Attribution . License Terms : Access for free at https://openstax.org/books/chemistry-2e/pages/1-introduction

All rights reserved content

  • What Is Electrolysis | Reactions | Chemistry | FuseSchool. Authored by : FuseSchool – Global Education. Located at : https://youtu.be/7uIIq_Ofzgw . License : Other . License Terms : Standard YouTube License

electrochemical cell in which electrolysis is used; electrochemical cell with negative cell potentials

process using electrical energy to cause a nonspontaneous process to occur

depositing a thin layer of one metal on top of a conducting surface

Chemistry Fundamentals Copyright © by Dr. Julie Donnelly, Dr. Nicole Lapeyrouse, and Dr. Matthew Rex is licensed under a Creative Commons Attribution-NonCommercial-ShareAlike 4.0 International License , except where otherwise noted.

Share This Book

  • Trending Categories

Data Structure

  • Selected Reading
  • UPSC IAS Exams Notes
  • Developer's Best Practices
  • Questions and Answers
  • Effective Resume Writing
  • HR Interview Questions
  • Computer Glossary

What are the Applications of Electrolysis?

What is electrolysis.

The process in which ionic substances are decomposed into simple substances by passing an electric current through them is known as electrolysis .

In other words, the process based on the fact that electrical energy can produce chemical changes is known as electrolysis .

Applications of Electrolysis

Nowadays the electrolytic process is widely used in various industrial applications. The major applications of the electrolysis are given below.

Extraction of Metal from their Ores

The electrolytic process is used for extracting out the pure metal from their ores, this process is known as electro-extraction . In the electro-extraction, the metal ore is treated with strong acid or is melted and then a DC current is passed through the resulting solution, the solution is decomposed and pure metal is deposited on the cathode.

Refining of Metals

Electrolysis is also used for refining of metals and the process is termed as electro-refining. In electro-refining , the anode of impure metal is placed in a suitable electrolytic solution. When DC current is passed through the solution, pure metal is deposited on the cathode.

Manufacturing of Chemicals

The electrolytic process is also used for manufacturing of various chemicals. When an electric current is passed through the solution of some compound, the compound gets breakdown into its constituent components which are liberated at the anode and cathode, which in turn can be collected.

Electro-Deposition

The electro-deposition is an electrolytic process, in which one metal is deposited over the other metal or non-metal. The electro-deposition is usually used for the decorative, protective and functional purposes.

Electroplating

An electrolytic process in which a metal is deposited over any metallic or non-metallic surface is called the electroplating . Electroplating is usually used to protect the metals from corrosion by atmospheric air and moisture.

Electro-deposition of Rubber

Electrolysis is also employed for electro-deposition of rubber . The rubber latex obtained from the tree consists of very fine colloidal particles of rubber suspended in water. These particles of rubber are negatively charged. On electrolysis of the solution, these rubber particles move towards the anode and deposit on it.

Electro-Metallization

The electrolytic process in which the metal is deposited on a conducting base for decorative and for protective purposes is termed as electro-metallization . Also, by using the electro-metallization process, any non-conductive base is made conductive by depositing a layer of graphite over it.

Electro-Facing

An electrolytic process in which a metallic surface is coated with a harder metal by electro-deposition in order to increase its durability is known as electro-facing .

Electro-Forming

Electrolysis is also used for electro-forming, it is the reproduction of an object by electro-deposition in order to increase its durability.

In the electro-forming, i.e. reproduction of medals, coins, etc., a mould is made by impressing the object in wax. The wax surface having exact impression of the object is coated by powdered graphite to make it conducting. This mould is then dipped in an electro-forming cell as cathode. After obtaining a coating of desired thickness, the article is removed and the wax core is melted out of the metal shell.

Electro-Typing

The electrotyping is an electrolytic process for forming metal parts that exactly reproduce a model. It is a special application of electro-forming and is mainly used to reproduce printing, set up tying, medals, etc.

The electrolysis process of deposition of an oxide film on a metal surface is known as anodizing . It is mainly used to increase the thickness of the natural oxide layer on the surface of the metal parts.

Electro-Polishing

The electro-polishing is an electrolytic process that removes materials from a metallic workpiece. It is also known as electrochemical polishing or electrolytic polishing.

Electro-polishing uses a combination of rectified current and a blended chemical electrolyte bath to remove flaws from the surface of a metal part.

Electro-Refining

Electro-refining is a method for purifying a metal using electrolysis. In the electro-refining process, the anode is made of impure metal and the impurities must be lost during the passage of metal from the anode to cathode during electrolysis.

Electro-Parting

An electrolytic process of separation of two or more metals is known as electro-parting or electro-stripping.

Electro-Cleaning

Electro-cleaning is the process of removing soil, scale or corrosion from a metallic surface. It is also known as electro-pickling . It is a form of electroplating which can be applied to all electrically conductive materials.

Manish Kumar Saini

Related Articles

  • What are the applications of DBMS?
  • What are the applications of OLAP?
  • What are the applications of clustering?
  • What are the applications of autoencoders?
  • What Are the Applications of Proteomics?
  • What are the Applications of IIoT?
  • What are the applications of C++ programming?
  • What are the Applications of Perl Programming?
  • What are the applications of Similarity Measures?
  • What are the applications of Association Rule?
  • What are the applications of Machine Learning?
  • What are the applications of Text Mining?
  • What are the applications of web mining?
  • What are the Applications of Pattern Mining?
  • What are the applications of Bipartite graphs?

Kickstart Your Career

Get certified by completing the course

Logo for UEN Digital Press with Pressbooks

72 Electrolysis

Learning objectives.

By the end of this section you will be able to:

  • Describe the process of electrolysis
  • Compare the operation of electrolytic cells with that of galvanic cells

Electrochemical cells in which spontaneous redox reactions take place ( galvanic cells ) have been the topic of discussion so far in this chapter. In these cells, electrical work is done by a redox system on its surroundings as electrons produced by the redox reaction are transferred through an external circuit. This final section of the chapter will address an alternative scenario in which an external circuit does work on a redox system by imposing a voltage sufficient to drive an otherwise nonspontaneous reaction, a process known as electrolysis . A familiar example of electrolysis is recharging a battery, which involves use of an external power source to drive the spontaneous (discharge) cell reaction in the reverse direction, restoring to some extent the composition of the half-cells and the voltage of the battery. Perhaps less familiar is the use of electrolysis in the refinement of metallic ores, the manufacture of commodity chemicals, and the electroplating of metallic coatings on various products (e.g., jewelry, utensils, auto parts). To illustrate the essential concepts of electrolysis, a few specific processes will be considered.

The Electrolysis of Molten Sodium Chloride

Metallic sodium, [latex]\text{Na}[/latex], and chlorine gas, [latex]\text{Cl}_2[/latex], are used in numerous applications, and their industrial production relies on the large-scale electrolysis of molten sodium chloride, [latex]\text{NaCl} (l)[/latex]. The industrial process typically uses a Downs cell similar to the simplified illustration shown in the figure below. The reactions associated with this process are:

The cell potential for the above process is negative, indicating the reaction as written (decomposition of liquid [latex]\text{NaCl}[/latex]) is not spontaneous. To force this reaction, a positive potential of magnitude greater than the negative cell potential must be applied to the cell.

write 3 application of electrolysis

The Electrolysis of Water

Water may be electrolytically decomposed in a cell similar to the one illustrated in the figure above. To improve electrical conductivity without introducing a different redox species, the hydrogen ion concentration of the water is typically increased by addition of a strong acid. The redox processes associated with this cell are

Again, the cell potential as written is negative, indicating a nonspontaneous cell reaction that must be driven by imposing a cell voltage greater than +1.229 V. Keep in mind that standard electrode potentials are used to inform thermodynamic predictions here, though the cell is not operating under standard state conditions. Therefore, at best, calculated cell potentials should be considered ballpark estimates.

write 3 application of electrolysis

The Electrolysis of Aqueous Sodium Chloride

When aqueous solutions of ionic compounds are electrolyzed, the anode and cathode half-reactions may involve the electrolysis of either water species [latex]\text{(H}_2\text{O, H}^+\text{, OH}^-)[/latex] or solute species (the cations and anions of the compound). As an example, the electrolysis of aqueous sodium chloride could involve either of these two anode reactions:

The standard electrode ( reduction ) potentials of these two half-reactions indicate water may be oxidized at a less negative/more positive potential (–1.229 V) than chloride ion (–1.358 V). Thermodynamics thus predicts that water would be more readily oxidized, though in practice it is observed that both water and chloride ion are oxidized under typical conditions, producing a mixture of oxygen and chlorine gas.

Turning attention to the cathode, the possibilities for reduction are:

Comparison of these standard half-reaction potentials suggests the reduction of hydrogen ion is thermodynamically favored. However, in a neutral aqueous sodium chloride solution, the concentration of hydrogen ion is far below the standard state value of 1 M (approximately 10 -7 M ), and so the observed cathode reaction is actually reduction of water. The net cell reaction in this case is then

This electrolysis reaction is part of the chlor-alkali process used by industry to produce chlorine and sodium hydroxide (lye).

CHEMISTRY IN EVERYDAY LIFE

Electroplating An important use for electrolytic cells is in electroplating . Electroplating results in a thin coating of one metal on top of a conducting surface. Reasons for electroplating include making the object more corrosion resistant, strengthening the surface, producing a more attractive finish, or for purifying metal. The metals commonly used in electroplating include cadmium, chromium, copper, gold, nickel, silver, and tin. Common consumer products include silver-plated or gold-plated tableware, chrome-plated automobile parts, and jewelry. The silver plating of eating utensils is used here to illustrate the process.

write 3 application of electrolysis

In the figure, the anode consists of a silver electrode, shown on the left. The cathode is located on the right and is the spoon, which is made from inexpensive metal. Both electrodes are immersed in a solution of silver nitrate. Applying a sufficient potential results in the oxidation of the silver anode [latex]\text{anode: Ag} (s) \rightarrow \text{Ag}^+ (aq) + \text{e}^-[/latex] and reduction of silver ion at the (spoon) cathode: [latex]\text{cathode: Ag}^+ (aq) + \text{e}^-\rightarrow \text{Ag} (s)[/latex] The net result is the transfer of silver metal from the anode to the cathode. Several experimental factors must be carefully controlled to obtain high-quality silver coatings, including the exact composition of the electrolyte solution, the cell voltage applied, and the rate of the electrolysis reaction (electrical current).

KEY TAKEAWAYS

Nonspontaneous redox processes may be forced to occur in electrochemical cells by the application of an appropriate potential using an external power source—a process known as electrolysis. Electrolysis is the basis for certain ore refining processes, the industrial production of many chemical commodities, and the electroplating of metal coatings on various products.

Chemistry End of Chapter Exercises

  • Write the half-reactions and cell reaction occurring during electrolysis of each molten salt below.  (a) [latex]\text{CaCl}_2[/latex]       (b) [latex]\text{LiH}[/latex]     (c) [latex]\text{AlCl}_3[/latex]       (d) [latex]\text{CrBr}_3[/latex]

This chapter is an adaptation of the chapter “ Electrolysis ” in Chemistry: Atoms First 2e by OpenStax and is licensed under a CC BY 4.0 license.

Access for free at https://openstax.org/books/chemistry-atoms-first-2e/pages/1-introduction

process using electrical energy to cause a nonspontaneous process to occur

element that conducts heat and electricity moderately well, and possesses some properties of metals and some properties of nonmetals

element that is shiny, malleable, good conductor of heat and electricity

Introductory Chemistry Copyright © by OpenStax is licensed under a Creative Commons Attribution-ShareAlike 4.0 International License , except where otherwise noted.

Share This Book

Library homepage

  • school Campus Bookshelves
  • menu_book Bookshelves
  • perm_media Learning Objects
  • login Login
  • how_to_reg Request Instructor Account
  • hub Instructor Commons
  • Download Page (PDF)
  • Download Full Book (PDF)
  • Periodic Table
  • Physics Constants
  • Scientific Calculator
  • Reference & Cite
  • Tools expand_more
  • Readability

selected template will load here

This action is not available.

Chemistry LibreTexts

13.3.4: Electrolysis- Using Electricity to Do Chemistry

  • Last updated
  • Save as PDF
  • Page ID 367895

    ↵

  So far, we have discussed how electricity can be produced from chemical reactions in batteries. Some reactions will, instead, use electricity to get a reaction to occur. In these reactions, electrical energy is given to the reactants, causing them to react to form the products. These reactions have many uses. For example, electrolysis is a process that involves forcing electricity through a liquid or solution to cause a reaction to occur. Electrolysis reactions will not run unless energy is put into the system from outside. In the case of electrolysis reactions, the energy is provided by the battery. Think of electrolysis and electrolytic cells as the opposite of electrochemical cells:

In an electrochemical cell, a spontaneous redox reaction is used to create an electric current; in an electrolytic cell, the reverse will occur—an electric current will be required in order to cause a non-spontaneous chemical reaction to occur. We will look at three examples of the electrolytic process, keeping our discussion on a very basic level—the electrolysis of molten sodium chloride, the electrolysis of water, and electroplating.

alt

If electrodes connected to battery terminals are placed in liquid sodium chloride, the sodium ions will migrate toward the negative electrode and be reduced while the chloride ions migrate toward the positive electrode and are oxidized. The processes that occur at the electrodes can be represented by what are called half-equations.

Reduction occurs at the positive electrode:

\[\ce{Na^+} + \ce{e^-} \rightarrow \ce{Na} \nonumber \]

Oxidation occurs at the negative electrode:

\[2 \ce{Cl^-} \rightarrow \ce{Cl_2} + 2 \ce{e^-} \nonumber \]

The overall reaction for this reaction is:

\[2 \ce{Na^+} + 2 \ce{Cl^-} \rightarrow 2 \ce{Na} + \ce{Cl_2} \nonumber \]

With appropriate treatment from the battery, it is possible to get the metal being reduced in an electrolysis process to adhere strongly to the electrode. The use of electrolysis to coat one material with a layer of metal is called electroplating . Usually, electroplating is used to cover a cheap metal with a layer of more expensive and more attractive metal. Many people buy jewelry that is plated in gold. Sometimes, electroplating is used to get a surface metal that is a better conductor of electricity. When you wish to have the surface properties of gold (attractive, corrosion resistant, or good conductor), but you don't want to have the great cost of making an entire object out of solid gold, the answer may be to use cheap metal to make the object and then electroplate a thin layer of gold on the surface.

alt

To silver plate an object like a spoon (plated silverware is less expensive than pure silver), the spoon is placed in the position of the cathode in an electrolysis set up with a solution of silver nitrate. When the current is turned on, the silver ions will migrate through the solution, touch the cathode (spoon) and adhere to it. With enough time and care, a layer of silver can be plated over the entire spoon. The anode for this operation would often be a large piece of silver, from which silver ions would be oxidized and these ions would enter the solution. This is a way of ensuring a steady supply of silver ions for the plating process.

  • Half-reaction at the cathode:

\[\ce{Ag^+} + \ce{e^-} \rightarrow \ce{Ag} \nonumber \]

  • Half-reaction at the anode:

\[\ce{Ag} \rightarrow \ce{Ag^+} + \ce{e^-} \nonumber \]

Some percentage of the gold and silver jewelry sold is electroplated. The connection points in electric switches are often gold plated to improve electrical conductivity, and most of the chromium pieces on automobiles are chromium plated.

Electrolysis of Molten Sodium Chloride

If we look at the latin roots of the word "electrolysis" we learn that it means, essentially, to "break apart" ( lysis ) using electricity. Our first example of an electrolytic cell will examine how an electric current can be used to break apart an ionic compound into its elements. The following equation represents the breaking apart of NaCl ( l ) :

2NaCl ( l ) → 2Na ( l ) + Cl 2 (g)

The half-reactions involved in this process are:

Notice that a negative voltage (-4.07V) results when we add up the half-reactions. This tells us that the overall reaction will NOT be spontaneous, and a minimum of 4.07 volts will be required for this reaction to occur.

As we shall see, our set-up will have a number of similarities to our electrochemical cells. We will need electrodes and an electrolyte to carry the electric current.

In our NaCl example, the electrodes will simply carry the current, but otherwise will not be directly involved in the reaction. The electrolyte will be the actual molten (melted) NaCl. The electrodes and electrolyte are both required to carry the electric current. Molten NaCl must be used because solid ionic compounds do not carry an electric charge.

Some key differences with an electrochemical cell set-up:

  • The two half-reactions are not separated by a salt bridge.
  • An electrochemical cell (or other source of electric current) will be required.

Other important items to note:

  • The anode of the electrolytic cell is the site of oxidation and the cathode is the site of reduction, just as in an electrochemical cell.
  • In an electrochemical cell, the anode is negative and cathode positive, but this is reversed in the electrolytic cell—the anode is positive and the cathode is negative.

Carefully study the diagram of our set up, taking special care to trace the path of the electrons. Unless electrons make a complete circuit, a reaction will not occur.

  • Electrons are "produced" in the battery at the anode, the site of oxidation.
  • The electrons leave the electrochemical cell through the external circuit.
  • These negative electrons create a negative electrode in the electrolytic cell, which attracts the positive Na + ions in the electrolyte. Na + ions combine with the free electrons and become reduced (2Na + + 2e - → Na).
  • Meanwhile, the negative Cl - become attracted to the positive electrode of the electrolytic cell. At this electrode, chlorine is oxidized, releasing electrons (Cl - → Cl 2 + 2 e - ).
  • These electrons travel through the external circuit, returning to the electrochemical cell.

Electrolysis of Water

Our second example of electrolysis and electrolytic cells involves the breakdown of water. We will find a situation very similar to the electrolysis of molten NaCl. The following equation represents the breaking apart of H 2 O ( l ) :

2H 2 O ( l ) → 2H 2 ( g ) + O 2 (g)

It may be more difficult to predict the half-reactions involved, but they are:

The set up will be very similar to our last example with some minor differences. Water does not carry a charge well, so an electrolyte is added to the water. Vinegar, a weak acid (acetic acid) may be used. To collect the hydrogen and oxygen gases produced, inverted test tubes are often added, as shown in our diagram below.

Again, take special care to trace the path of the electrons. Unless electrons make a complete circuit, a reaction will not occur.

2H 2 O ( l ) + 2e - → H 2( g ) + 2 OH - (aq)

2H 2 O ( l ) → O 2 (g) + 4H + (aq) + 4e -

  • These negative electrons create a negative electrode in the electrolytic cell which causes the reduction of water. Note that the area around this electrode will become basic as OH - ions are produced.
  • Meanwhile, the positive electrode water will undergo oxidation.
  • Electrons produced during this oxidation process will return to the electrochemical cell.

A note about the balanced equation for the electrolysis of water:

You may notice from the half reactions that adding up the equations doesn't initially give us our net equation of:

Once you balance for electrons (multiply the reduction equation by 2), you'll find that the equations actually add up to:

6H 2 O ( l ) → 2H 2 ( g ) + O 2 (g) + 4H + (aq) + 4 OH - (aq)

The hydrogen and hydroxide ions will combine to form 4 moles of H 2 O ( l ) . Finding our net amount of H 2 O ( l ) involved gives us our final equation:

  • Electrochemical cells are composed of an anode and cathode in two separate solutions. These solutions are connected by a salt bridge and a conductive wire.
  • An electric current consists of a flow of charged particles.
  • The electrode where oxidation occurs is called the anode and the electrode where reduction occurs is called the cathode.
  • In electroplating, the object to be plated is made the cathode.
  • Electrochemical cell - An arrangement of electrodes and ionic solutions in which a redox reaction is used to make electricity (also known as a battery).
  • Electrolysis - A chemical reaction brought about by an electric current.
  • Electroplating - A process in which electrolysis is used as a means of coating an object with a layer of metal.
  • Current Electricity
  • Electrolysis And Electroplating

Electrolysis and Electroplating

We have heard of the coating of one metal on another metal using the process of electrolysis. We have also heard about copper plating and silver plating in our daily lives. In this article, let us discuss in brief electrolysis and electroplating along with their applications.

What Is Electrolysis?

The word “electrolysis” was introduced by Michael Faraday in the 19th century. In chemistry, electrolysis is a method that uses a direct current (DC) to drive a non-spontaneous chemical reaction. This technique is commercially significant as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell.

We can define electrolysis as:

The process by which ionic substances are decomposed into simpler substances when an electric current is passed through them.

Electrolysis Process

The fundamental process of electrolysis is the interchanging of ions and atoms by the addition or removal of electrons from the external circuit.

Electrolysis Process

  • Ionic compounds contain charged particles called ions. For example, sodium chloride contains positively charged sodium ions and negatively charged chlorine ions.
  • The ions must be free to move in order to start the electrolysis process. When an ionic substance is dissolved in water or melted then the ions are free to move. During electrolysis, positively charged ions move to the negative electrode and negatively charged ions move to the positive electrode. Then positively charged ions receive electrons and negatively charged ions lose electrons. Both the products of the dissociation get collected at the electrodes.
  • For instance, if electricity is passed through molten sodium chloride, the sodium chloride is broken into sodium and chlorine, and they collect at their respective electrodes. The metals get precipitated and the gases escape. This ability to break down a substance with a current is used in many ways.
  • Electrolysis is widely used for electroplating.

Why Is Aluminium No Longer a Precious Metal?

In medieval times, aluminium was almost as precious as silver. Its extraction was not known and pure aluminium was extremely hard to come by. The application of electrolysis turned aluminium from a precious metal to one of the most largely used metals by humans. Aluminium is extracted from its ore by the process of electrolytic decomposition and deposition.

Uses of Electrolysis

  • Electrolysis is done for coating one metal on another.
  • The industrial use includes various metals such as aluminium, magnesium, chlorine, and fluorine etc.

Watch and Learn about the Basic Concepts of Current Flow

write 3 application of electrolysis

What Is Electroplating?

Electroplating is a process that uses an electric current to reduce dissolved metal cations so that they form a thin coherent metal coating on an electrode.

  • Electroplating is a process that uses electric current to reduce dissolved metal ions by the use of electrolysis, to obtain the dissolved metal ions at the other electrode, mostly in the form of a uniform coating.
  • It is the process of plating one metal onto another by hydrolysis, most commonly for decorative purposes or to prevent corrosion of metals.
  • There are also specific types of electroplating such as copper plating, silver plating, and gold plating. It allows the manufacturers to make the product with economical materials and then coat the metals to add properties such as rust proofing, improving its appearance and improving its strength.

How Does Electroplating Work?

The positive electrode should be the metal that you want to coat the object with. The negative electrode will be the object that should receive the deposit of the electroplating metal. The electrolyte should be a solution of the coating metal, as it is a metal nitrate or sulphate. We can examine electroplating better with a few examples;

Silver Plating:

Silver Plating

The object that is to be plated is connected to the negative terminal of the power supply. A piece of silver is connected to the positive terminal. The electrolyte is a silver nitrate solution. This process can also be used to purify silver. We can take an impure block at the positive terminal and a strip of pure silver to the negative. The pure silver is deposited via electrolysis at the negative terminal, leaving behind the impurities.

Copper Plating:

Copper Plating

The object that is to be plated is connected to the negative terminal of the power supply. A piece of copper is connected to the positive terminal. The electrolyte is a copper Sulphate solution.

Like with silver plating, copper deposition can be used to purify copper. In this case, both the electrodes are made from copper. As the impure positive electrode gradually disappears leaving the impurities behind, the negative electrode gradually gets coated with pure copper.

Read More: Corrosion of Metals

Difference between Electrolysis and Electroplating

Following is the table explaining the difference between electrolysis and electroplating:

Electrolysis of Water

Electrolysis of water is the decomposition of water into oxygen and hydrogen gas. This is achieved by passing an electric current through the water. Two electrodes are placed in the container with water. The electric current is passed through these electrodes. Stainless steel or platinum are used for the making of the electrodes.

The negatively charged electrode that is cathode is the place where the hydrogen will get accumulated while the positively charged electrode that is anode is the place where the oxygen will get accumulated. Electrolysis of water is also an example of decomposition which is also known as electrolytic decomposition.

Frequently Asked Questions – FAQs

What is the principle behind electroplating.

Hydrolysis is the principle behind electroplating. For electroplating to take place, metal salt should be present in the aqueous solution. The end product of electroplating is water molecule. Thus, hydrolysis is the principle behind electroplating.

Mention the purposes of electroplating.

Following are the main purposes of electroplating:

  • For protecting metals against corrosion
  • To decorate the articles by giving them shiny appearances
  • To fix the worn out materials

Define nickel strike.

Nickel strike is also known as flash layer and is defined as the process in which electric current is used to coat a conductive material along with a thin layer of nickel.

How many ions can be discharged during electrolysis?

During electrolysis, 2 out of 4 ions can be discharged during electrolysis.

Name the electrolyte used in lead-acid cells.

Dilute H 2 SO 4 is the electrolyte that is used in lead-acid cells.

Name the metal that is purified using electrolysis.

Copper is the metal that is purified using electrolysis.

Hope you have learned about electrolysis and electroplating in a detailed and interesting way. To know more about various science and maths concepts. Visit BYJU’S – The Learning App.

Quiz Image

Put your understanding of this concept to test by answering a few MCQs. Click ‘Start Quiz’ to begin!

Select the correct answer and click on the “Finish” button Check your score and answers at the end of the quiz

Visit BYJU’S for all Physics related queries and study materials

Your result is as below

Request OTP on Voice Call

Leave a Comment Cancel reply

Your Mobile number and Email id will not be published. Required fields are marked *

Post My Comment

write 3 application of electrolysis

The way they explain is great. They make hard concepts easy with their explanation. 👌👏👍

I like this topic seriously

The explaining is very good I like it 👏👏👍

write 3 application of electrolysis

  • Share Share

Register with BYJU'S & Download Free PDFs

Register with byju's & watch live videos.

close

IMAGES

  1. Electrolysis Process Diagram

    write 3 application of electrolysis

  2. Applications of electrolysis

    write 3 application of electrolysis

  3. Understand How Electrolysis Works Worksheet

    write 3 application of electrolysis

  4. IGCSE Chemistry-Electrolysis practice

    write 3 application of electrolysis

  5. Electrolysis and Electroplating

    write 3 application of electrolysis

  6. Electrolysis: Definition, Mechanism, Meaning & Application

    write 3 application of electrolysis

VIDEO

  1. KraziEduMusic Application of electrolysis

  2. ELECTROCHEMISTRY:APPLICATION-ELECTROPLATING AND ELECTROLYSIS

  3. Can electrolysis be made more efficient?

  4. Electroplating |Application of electrolysis |electrochemistry |polytechnice1st y.|MSBTE| Diploma F.Y

  5. GCSE

  6. Explain the term electrolysis. Write the reactions at cathode ad anode when following s

COMMENTS

  1. Electrolysis: Definition, Process, Equations, Examples, and Applications

    Examples of Electrolysis [1-4] 1. Water (H 2 O) Water can undergo electrolysis in the presence of an electrolyte like acid or base. The presence of acid improves the electrical conductivity by increasing the hydrogen ion (H +) concentration. Examples of such electrolytes are sulfuric acid (H 2 SO 4) and salt of sodium nitrate (NaNO 3 ).

  2. Definition, Process, Applications, Electrolysis of Water

    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:

  3. Applications of Electrolysis: Top 12 Applications

    The following points highlight the top twelve applications of electrolysis. The applications are: 1. Extraction of Metals 2. Refining of Metals 3. Production of Chemicals 4. Electroplating 5. Deposition of Alloys 6. Electroforming 7. Electrotyping 8. Electrofacing 9. Electrometallisation 10. Electrodeposition of Rubber 11. Anodizing 12. Electropolishing. Application # 1. Extraction of Metals ...

  4. Electrolysis

    electrolysis, process by which electric current is passed through a substance to effect a chemical change. The chemical change is one in which the substance loses or gains an electron (oxidation or reduction). The process is carried out in an electrolytic cell, an apparatus consisting of positive and negative electrodes held apart and dipped into a solution containing positively and negatively ...

  5. 17.7: Electrolysis

    The Electrolysis of Molten Sodium Chloride. Metallic sodium, Na, and chlorine gas, Cl 2, are used in numerous applications, and their industrial production relies on the large-scale electrolysis of molten sodium chloride, NaCl(l).The industrial process typically uses a Downs cell similar to the simplified illustration shown in Figure \(\PageIndex{1}\).

  6. 14.5: Electrolysis

    Figure 14.5.1 14.5. 1: Electrolysis. In an electrolytic cell, electricity is forced through the cell to induce a non-spontaneous redox reaction. Here, the redox reaction 2H 2 O → 2H 2 + O 2 is being caused by the introduction of electricity, which is supplied by the battery. Electrolysis has many applications.

  7. Electrolysis

    Electrolysis also has numerous other applications. For example, fluorine cannot be prepared by any spontaneous chemical reaction. It was found in 1886 by the French chemist Henri Moissan that passing electrical current through a molten mixture of potassium fluoride and hydrogen fluoride produces Fluorine, and it is still prepared commercially ...

  8. Electrolysis

    In an electrolytic cell, electricity is forced through the cell to induce a nonspontaneous redox reaction. Here, the redox reaction 2 H 2 O → 2 H 2 + O 2 is being caused by the introduction of electricity, which is supplied by the battery. Electrolysis has many applications. For example, if NaCl is melted at about 800°C in an electrolytic ...

  9. Electrolysis

    Figure 14.5 Electrolysis. In an electrolytic cell, electricity is forced through the cell to induce a nonspontaneous redox reaction. Here, the redox reaction 2 H 2 O → 2 H 2 + O 2 is being caused by the introduction of electricity, which is supplied by the battery. Electrolysis has many applications. For example, if NaCl is melted at about ...

  10. Applications of Electrolysis

    Applications of Electrolysis . Download PDF for free. Application of Electrolysis - example. Electrolysis is applied in the various field. It is required for extraction and purification of metals. ... Write down the uses/application of electrolysis. View Answer. Electrolysis rules of Faraday's states that mass depends on electrodes is ...

  11. Give three applications of electrolysis.

    Solution. The three applications of electrolysis are: 1. Electroplating - This process applies the method of electrolysis in plating a thin layer of superior metal (silver, gold, nickel, chromium, etc.) over the inferior metal. This is done to prevent the inferior metal (iron, copper, etc.) from corrosion and also for improving the look of the ...

  12. Industrial Applications of Electrolysis

    For example, electrolysis of sodium chloride solution, also known as brine, can produce sodium hydroxide, hydrogen, and chlorine. These chemicals have a variety of uses in manufacturing and other applications. Another example of this application is the electrolysis of water to produce hydrogen and oxygen. Electroplating

  13. Top 7 Applications of Electrolysis

    The main practical applications of electrolysis are: 1. Extraction of Metals 2. Refining of Metals 3. Production of Chemicals 4. Electroplating 5. Electro-Typing 6. Electro-Forming 7. Electro-Cleaning. Application # 1. Extraction of Metals: There are two methods of extraction of metal on the basis of physical states of the ore. In the first method the ore is treated with a strong acid to ...

  14. Applications of Electrolysis for IIT JEE

    Electrolysis is a process of chemical decomposition produced by passing an electric current through a liquid or solution containing ions. The key process of electrolysis is the exchange of atoms and ions from the external circuit by removing or adding electrons. Read more about the Applications of Electrolysis for IIT JEE (Main and Advanced) at Vedantu.com.

  15. 1.10: Electrolysis- Using Electricity to Do Chemistry

    Electroplating Figure 16.7.1: An electrical current is passed through water, splitting the water into hydrogen and oxygen gases. If electrodes connected to battery terminals are placed in liquid sodium chloride, the sodium ions will migrate toward the negative electrode and be reduced while the chloride ions migrate toward the positive electrode and are oxidized.

  16. Electrolysis

    In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction.Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell.The voltage that is needed for electrolysis to occur is called the decomposition potential.

  17. 17.6 Electrolysis

    The Electrolysis of Water. Water may be electrolytically decomposed in a cell similar to the one illustrated in Figure 17.6.2. To improve electrical conductivity without introducing a different redox species, the hydrogen ion concentration of the water is typically increased by addition of a strong acid.

  18. What are the Applications of Electrolysis?

    Electrolysis is also used for electro-forming, it is the reproduction of an object by electro-deposition in order to increase its durability. In the electro-forming, i.e. reproduction of medals, coins, etc., a mould is made by impressing the object in wax. The wax surface having exact impression of the object is coated by powdered graphite to ...

  19. Electrolysis

    Nonspontaneous redox processes may be forced to occur in electrochemical cells by the application of an appropriate potential using an external power source—a process known as electrolysis. Electrolysis is the basis for certain ore refining processes, the industrial production of many chemical commodities, and the electroplating of metal ...

  20. Explain the applications of electrolysis

    Electrolysis is the process by which electric current is passed through a substance to effect a chemical change. The chemical change is one in which the substance loses or gains an electron (oxidation or reduction). The process is carried out in an electrolytic cell, an apparatus consisting of positive and negative electrodes held apart and ...

  21. 13.3.4: Electrolysis- Using Electricity to Do Chemistry

    This page titled 13.3.4: Electrolysis- Using Electricity to Do Chemistry is shared under a mixed license and was authored, remixed, and/or curated by Anonymous. Galvanic cells produce electricity from chemical reactions. Some reactions will, instead, use electricity to get a reaction to occur. In these reactions, electrical energy is given to ...

  22. Electrolysis and Electroplating

    Electrolysis. Electroplating. Electrolysis refers to the breaking apart of a molecule by the means of the electrochemical reaction. Electroplating refers to the passage of current through the solution with metal such that it gets deposited on one of the electrodes. Electrolysis is good for carrying out the non-spontaneous chemical reactions.

  23. Applications of Electrolysis Electroplating Electroforming

    During electrolysis process, the electrolyte metal will be deposited on the graphite coated impressed surface of the mold. After obtaining a layer of desired thickness, the article is removed and the wax is melted to get the reproduced object in form of metal shell. A popular use of electroforming is reproduction of gramophone record dices.