Faraday's Laws Of Electrolysis

8.3 Faraday's Laws Of Electrolysis

Michael Faraday in 1833 discovered the quantitative laws governing the process of electrolysis, which are known as Faraday's laws of electrolysis. There are two laws of electrolysis put forward by Faraday.

1. Faraday's First Law of Electrolysis:

It states that the amount of any substance deposited or liberated at an electrode during electrolysis is directly proportional to the-quantity of current passed through the electrolyte.


If 'w' is the weight or amount of a substance deposited or liberated and 'A' ampere of current is passed for t seconds, then according to the law:

W    ∞    A x t

Or   w    =    Z A t


Where Z is a constant, known as "electro chemical equivalent" of the substance. If one ampere of current is passed for one second; then w = Z. This means when one ampere of current is passed for one second; then the weight or amount of the substance deposited or liberated is exactly equal to its electrochemical equivalent. The current in ampere, multiplied by the time in second is known as coulomb which is the electric charge.

Ampere (A) x time (s)    = coulomb (C)

Electrochemical Equivalent:

Electrochemical equivalent of a substance may be defined as the weight (amount) of the substance deposited or liberated, when one coulomb of electric charge is passed through an electrolyte. It is denoted by Z and in S.I unit it is expressed in Kg / coulomb. Each element has its own electrochemical equivalent.

Example - 1

A current of 0.5 ampere was passed through a solution of CuS04 for' one hour. Calculate the mass of copper metal deposited on the cathode. Electrochemical equivalent of Cu=0.000329g/C = 3.29 X 10-4 g/C

Or        3.294X10-7 Kg/C.



1.       Current in ampere (A) = 0.5

2.       Time in second (1 hr) = 1 x 60 x 60        = 3600 s

3.       Z for Cu metal         = 3.294x10-4 g/C = 3.294x10-7Kg/C


w   =    Zx Axt

     =    3.294 X 10-7 x 0.5 x 3600 =    5.929 x 10-4 Kg

     =    5929.2 x 10-4Kg

    Mass of copper metal deposited g = 5.929 x 10-4 Kg or 0.5929

Example - 2

A current of 10 amperes was passed for 15 minutes in a solution of silver nitrate (AgNO3). The mass of silver deposited was found to be 1.0062 x 10-2 Kg. Calculate the electrochemical equivalent (Z) of Ag metal.



1. Current in ampere= 10

2. Time in seconds (15 minutes) = 15 x 60=900 s

3. Mass of Ag metal deposited (w) = 1.0062 x 10-2 Kg


          w = Z A t

Or Z= w/At = 1.0062x10-2Kg/10Ax900s

=1.0062x10-2/9x10-3 = 1.0062x10-2x10-3/9

 = 0.1118x10-5Kg/C

= 1.118xl0-6 Kg/C

= 0.00118   g/C.

2. Faraday's Second Law Of Electrolysis:

It states that the masses of different substances deposited or liberated, when same quantity of current is passed through different electrolytes, connected in series are proportional to their chemical equivalent masses.

Consider three different electrolytes, AgNO3, CuSO4 and Al (N03)3 solutions, connected in series. Same quantity of current is passed through them, then the masses of Ag, Cu and Al, deposited on their respective electrodes would be directly proportional to their equivalent masses. Atomic mass of the element

According to Faraday if exactly 96,500 coulombs of electric charge is passed then the mass of Ag deposited would be equal to 108g (108/1), that of copper is 31.75g (63.5/2) and Al is 9g (27/3) which are their equivalent masses respectively.

                Equivalent mass of an element =Atomic mass of an element

                                                                Valency of the element

The current of 96,500 coulombs is called as one Faraday (F) charge after the name of Faraday. Thus Faraday is defined as the quantity of charge which deposits or liberates exactly one gram equivalent of a substance.

Relationship between Equivalent Mass and Electrochemical Equivalent:

Since 96,500 C (IF) electric charge is required to liberate one gram equivalent mass of a substance, so it is clear that the gram equivalent mass of a substance is 96,500 times greater than its electrochemical equivalent.

         Gram Eq.mass = 96,500 x E.C.E (Z)

In other words if e is the gram equivalent mass and Z is the E.C.E, then we can write it as:

e = 96,500 x Z or   e = F x Z.


1. Ampere:

It is the basic unit of current in the international system (S.I). It may be defined as the current when passed through a circuit for one second, can liberate 0.001118 g or 1.118x 10-6 Kg of Ag from silver nitrate solution.

2. Coulomb:

It is the SI unit of electric charge and is defined as the quantity of charge when one ampere of current is passed for one second.

C    =     ampere x time (s) 1 coulomb   

      =     1 A      x 1 s

A coulomb is equivalent to ampere multiplied by second, for example, when a current of 0.5 ampere flowing for 80 seconds, gives a charge 0.5 x 80 = 40 coulombs.

If we are given the current and the time of electrolysis, we can calculate the amount of the substance produced, at an electrode. On the other hand if we are given the amount of the substance produced at an electrode and the length of the time of electrolysis, we can determine the current or electrical charge.


When an aqueous solution of copper sulphate is electrolysed, copper metal is deposited at the cathode.

Cu2+ (aq) + 2e- ----- > Cu(s) -- at cathode

If a constant current was passed for 5 hours and 404 mg of Cu was deposited. Calculate the current passed through CuSO4-.


Amount of Cu deposited = 404mg = 0.404g.

Gain of 2e- means 2F electric charge

Atomic mass of Cu = 63.5

According to cathode reaction

63.5 g of Cu is deposited by 2 F electric charge

 0.404g of Cu is deposited by 2/63.5 x 0.404 = 0.0127 F

We know

IF = 96,500 coulomb

0.0127F = 0.0127 X96, 500 = 1225.6C

Coulomb = Ampere x t sec. (time = 5hours)

Ampere= C/ts = 1225.6/5x60x60 = 0.0680

= 6.80 x 10-2 ampere


How many grams of oxygen gas is liberated by the electrolysis of water after passing 0.0565 ampere for 185 second.


      Equation:     2H2O ----- > O2 (g) + 4H+ + 4e-

According to equation (four) Faraday is required to liberate 32 g of 62


Current in ampere        = 0.0565

Time in second            = 185

Coulomb                     = ampere x time (s)

                                  = 0.0565x185 - 10.45 C.

                               F =C/96,500 = 10.45/96,500 = 0.000108 F



4 F electric charge liberates 32 g O2


0.000108 Electric charges liberates 32/4 x 0.000108

=    0.000864 g =    8.64 xl0-4g O2

8.4 Uses Of Electrolysis

Electrolysis is an important process, used for the extraction of certain metals, from their ores. It is also used in electroplating. Electroplating means to coat one metal on other metal through the process of electrolysis in order to protect the baser metals from corrosion and to make them more attractive.

In 1852, the price of Aluminum metal was very high and Al metal was considered as an expensive oddity, because of its slivery - white shining property but by 1890, the price of Al - metal had fallen very low, a fraction of the price of silver. What had happened?

Aluminum metal is a reactive metal which makes it difficult to extract it from its ore. Before 1886 the only way to get it, was by heating its salt Aluminum chloride with sodium metal. Sodium itself was expensive which made it even more expensive.

AlCl3(s) + 3Na(s) heat> Al(s) + 3NaCl(s)

Then in 1886 a new process was developed which involved the electrolysing a molten Aluminum compound. The same basic method of electrolysis is still used for the extraction of Al-metal. Now-a-days Al is extracted by the electrolysis of Alumina (A1203) which is obtained from the chief ore of Al, bauxite (A1203 nH2OJ. Due the process of electrolysis, the price of Aluminum quickly falls. Today Aluminum is so cheap that if you use disposable plates of Aluminum for eating and throw them away after words.

Many metals are extracted by the electrolysis of their molten compounds, usually their chlorides, because the chlorides of most of the metals have low melting points. For example sodium metals is extracted by the electrolysis of molten sodium chloride to deposit Na metal at cathode by Down's process. Many metals are purified into pure metals by the process of electrolysis for example impure copper (blister Cu) is purified by the process of electrolysis. In this process, the impure copper i.e. Blister copper is made as anode in the electrolytic cell, cathode is a thin plate or sheet of pure copper metal and the two electrodes are dipped in the electrolytic solution of copper sulphate (CuSO4), containing few drops of sulphuric acid. The two electrodes are connected with a battery (source of electricity), when an electric current is passed through the electrolytic solution. The copper anode dissolves in the solution as Cu2+ ions which move towards cathode and gain electrons to get neutralized, depositing pure copper metal over cathode plate. Most of the impurities of anode fall to the bottom of the cell, called as "anode mud". Copper thus deposited at cathode is 99.99% pure. In this way copper anode (Blister copper) dissolves completely to form pure copper at cathode. This process of electrolysis is similar to electroplating.


Electroplating is the process of electrolysis which is used to coat one metal onto another. Usually the object to be electroplated is made up of cheaper or baser metal, such as iron, steel etc. It is then coated with a thin layer of more attractive corrosion - resistant and costly metal, like silver, gold, nickel, chromium, tin etc. The cost of the finished product is far less than the objects entirely made of these metals. Gold coated object is much cheaper than the gold object.

1. Nickel Plating:

A cell for electroplating of nickel is shown in the fig 8.4. A piece of pure nickel is the anode and the spoon or any object to be nickel plated is cathode. A solution of nickel sulphate (NiSO4) is used as the electrolyte in the electrolytic cell. The two electrodes are joined with a battery (an external source of direct electric current). On passing the electric current, the anode which is Ni, dissolves in the electrolytic solution forming Ni2+ions by the loss of electrons. Ni2+ions from the solution move towards the cathode, where they gain electrons and are reduced to Ni metal on the surface of spoon (cathode)

Ni(s) (anode) ----- > Nia2+ (aq) + 2e-     (Anode reaction oxidation)

Ni2+ (aq) + 2e- ----- >   Ni(s)         (Cathode reaction reduction)

The net reaction is simply the transfer of Ni as Ni2+ through NiSO4 solution towards the cathode i.e. spoon and get it coated with Ni metal on the surface. The sum of reduction and oxidation is: 

2. Chromium Plating:

Chromium metal can also be coated over baser or cheap metals by the process of electrolysis i.e. electroplating.

A cell of chromium electroplating is shown in fig. 8.5. A piece of chromium metal is the anode and the spoon or any other object to be chromium plated is the cathode. A solution of chromium sulphate [Cr2 (SO4)3J is used as an electrolyte in the electrolytic cell. The two electrodes are joined with a battery (an external source of direct electric current). On passing electric current, the anode which is chromium dissolves in the solution, forming Cr3+ ions by the loss of electrons, Cr3+ ions from the solution move towards the to cathode, where they gain electrons and are reduced to deposit chromium metal on the surface of cathode.

Cr(s) (anode) ----- > Cr3+ (aq) + 3e-     (Anode reaction oxidation)

Cr3+ (aq) 3e- ----- > Cr(s)   (Cathode reaction reduction)


The net process is simply the transfer of Cr as Cr3- ions through Cr2 (SO4)3 solution towards cathode and coated it with Cr- metal. The sum of oxidation and reduction is:

Looking for Electro plate:

Look around in your home for articles that are electroplated. Decide as far as you can :

1.    What metal is plated on top.

2.    What metal is underneath.

3.    What is the purpose of the plating?

Choose one of the items and look very closely at it, specially look at cooking oil container. Is there any place where the plating is less good?

8.5 Electrochemical Cells:

The cell which is used to convert chemical energy into electrical energy or vice verse is called electrochemical cell. An electrochemical cell which converts chemical energy into electrical energy is known as Galvanic or voltaic cell. This is a very strange device to produce electric current.

Take strips of zinc and copper metals, sticking them into a lemon, producing electricity. How this happens? The two dissimilar metal strips (Zn and Cu) and the electrolyte of lemon juice are the key components of the device that convert. Chemical energy into electrical energy, you try it of your own. Insert the Zn and Cu strips into a lemon and join the two strips by a wire outside, you will feel the electricity passing through wire joining the two strips.

The simplest of the Galvanic or Voltaic cells is Daniell cell.

Daniell cell:

A Daniell cell is the simplest of the Galvanic or Voltaic cells which is used to convert chemical energy into electrical energy spontaneously. Daniell cell consists of two half cells. One half cell is Zinc rod (Zn - metal) dipped in IM ZnSO4 solution and the other half cell is Copper rod (Cu-metal) dipped in IM CuSO4 solution. The two half cells or single electrodes  are connected together to form a complete cell. The two half cells should be separated from each other by a porous partition or a salt bridge, when both the electrodes are connected externally through a voltmeter by means of metal wire. The cell starts producing electric current at once. Zn undergoes oxidation to form Zn2+ ions by the loss of 2 - electrons to go into ZnSO4 solution. Zn acts as anode or negative electrode. The electrons which are free at Zn - electrode travel through the wire externally to Cu - electrode. These electrons are accepted by Cu2+ions of CuSO4 solution and Cu2+ions undergo reduction to deposit copper metal at Cu - electrode which acts as cathode or positive electrode. In this process Zn - electrode dissolves in the solution of ZnSO4 and reduces in size, while copper electrode grows in size due to the deposition of Cu - metal.


Heart Beats and Electrocardiography (ECG)

The human heart, generally in a whole day pumps more than 7000 liters of blood through the circulatory system. We generally think of the heart as a mechanical device; a muscle that circulates blood via regularly spaced muscular contractions. However, in the late 1800 two pioneers in electricity Luigi Galvani and Alesandro Volta discovered that the contractions of heart are controlled by electrical phenomenon, as are nerve impulses throughout the body. These electrical impulses that cause the contraction of heart muscle are strong enough to be detected at the surface of the body. This observation formed the basis for electrocardiography (ECO). It is quote striking that although the hearts major function is the mechanical pumping of blood, it is most easily monitored by using the electrical impulses generated by voltaic Cell.


Cell Reaction

At Zn - Electrode (Anode)

Zn(s) ----- > Zn2+ (aq) + 2e-

At Cu - Electrode (Cathode)

Cu2+ (aq)   + 2e- ----- > Cu(s)

The total cell reaction is the sum of two half cell reactions.

Zn(s) + Cu2+ (aq) ----- > Zn2+ (aq) + Cu(s)

The function of salt bridge or porous partition is to prevent the mixing of two solutions (ZnSO4 and CuSO4) and allows the ions to move through from one part to another. Zn2+ ions from the anode compartment move into the cathode compartment, while negative ions: SO42- ions move from cathode compartment to anode compartment through the porous partition or salt bridge. It maintains the electrical neutrality in the two electrolytic solutions

The cell voltage in Daniell cell is 1.10 volt.

8.6 Batteries:

In every day life, we use the devices to produce electricity by the chemical reactions, known as batteries. A flash light "battery" consists of single voltaic cell with two electrodes in contact with one or more electrolytes, some times a distinction is made between the terms "cell" and "battery". A battery is an assembly of two or more voltaic cells, connected together in series. By this definition automobile or motor battery is a true battery. The most common types of cells or batteries are described as follows:

1. Dry Cell:

It is a primary cell, which is used to convert chemical energy into electrical energy. It is used in most of the flash lights, calculators, clocks, transistors and in portable electronic devices. It is an irreversible cell. In a dry cell (fig 8.7), there is an outer zinc vessel which acts as anode and inert carbon (graphite) rod which acts as cathode. The graphite rod is surrounded by a mixture of mangnese dioxide (MnO2) and carbon powder. The electrolyte is a moist paste of ammonium chloride (NH4Cl) and zinc chloride (ZnCl2). The cell is called a dry cell because there is no free flowing liquid. The concentrated electrolytic solution is thickened into a gel like paste by an agent such as starch. The upper top position of the cell is sealed with wax (sealing material). A copper cap is fitted on the top of carbon rod (cathode) to make the electrical contact. The whole cell is covered with a safety cover.

   The cell Diagram is as:

When zinc and graphite electrodes are connected by a metallic wire, Zn gets oxidized to form Zn2+ions which pass into the wet paste leaving behind electrons on the Zn container and the electrons move from Zn electrode to carbon electrode through the external circuit. The cell reactions are complex.

2. Lead-Storage Battery (Motor - Battery)

Lead storage battery is used in automobiles. It is a secondary battery and is a reversible cell which can be restored to its original condition. The battery can be used through repeated cycles of discharging and recharging.

 Fig 8.8 shows a portion of the cell in the lead storage battery. There are several anodes and several cathodes which are connected together in series; about six cells are connected together. Each cell has a voltage of 2V and overall voltage when six cells are connected together in series would be 12V.

In lead storage battery anodes are the lead alloy and cathodes are made up of red lead dioxide (PbO2). The electrolyte is dilute sulphuric acid. Which having concentration of 30%

As the cell reaction precedes PbSO4(s) precipitates and partially coats both the electrodes, the water formed dilutes the sulphuric acid. The battery is said to be discharged. Now by connecting the battery to an external electrical source, we can force the electrons to flow in the opposite direction, i.e. the net cell reaction can be reversed and the battery is recharged.

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