Chapter 3: Electrochemistry


Redox Reactions

Last Updated : 18 Oct, 2023

Redox Reactions are oxidation and reduction reactions that happen simultaneously in a chemical reaction and in this, the reactant undergoes a change in its oxidation state. Redox stands for Reduction – Oxidation. Redox reaction is a common term used in both Chemistry and Biology. They are a certain type of chemical reaction in which the substrate’s oxidation states change.

A reduction is a decrease in the oxidation state or a gain in electrons, whereas oxidation is the loss of electrons or an increase in the oxidation state. A redox reaction involves the transfer of electrons between two species. Redox reactions can occur in two ways, by the electron transfer and by the atom transfer of the chemical species involved in a chemical reaction. Redox Reaction is important for class 10 and class 11 students.

In this article, you will read about what is a redox reaction, its definition, types, and examples of redox reactions.

Table of Content

What is a Redox Reaction?

Redox reaction in Chemistry involves the alteration of oxidation states of the atoms. These reactions involve the actual shifting or transfer of electrons of the chemical species involved in the reaction. So, in this reaction, one species loses electrons while the other gains electrons.

Redox reaction can be determined by the change or difference in the oxidation state of two atoms. If there is no change or difference in oxidation number then there is no redox reaction can take place. 

Redox Reaction Definition

A redox reaction, short for reduction-oxidation reaction, is a chemical process in which the oxidation states (or oxidation numbers) of one or more substances involved in the reaction change.

Types of Redox Reactions

There are four different types of redox reactions, that are:

  • Decomposition Reaction
  • Combination Reaction
  • Displacement Reaction
  • Disproportionation Reaction

Decomposition Reaction

A chemical is broken down into distinct components in this type of reaction. It is the inverse reaction of a combination reaction. In a displacement reaction, for example, the atom gets replaced by an atom of another element. 

A chemical equation will be used to depict the chemical reaction. It denotes the transition from reactants to products. The reactant side is represented on the left, and the result of the reaction is represented on the right.

For examples,

2NaH → 2Na + H2

2H2O → 2H2 + O2

There are three types of decomposition reactions:

  1. Thermolysis – Thermolysis is heat-induced decomposition.
  2. Electrolysis is the decomposition of matter caused by electricity.
  3. Photolysis – Photolysis is decomposition due to light.

Combination Reaction

These reactions, which are the inverse of decomposition processes, require the combination of two chemicals to generate a single compound in the form of A + B → AB.

The outcome of a combination reaction between a metal and a nonmetal is an ionic solid. As an example, lithium can react with sulfur to form lithium sulfide. When magnesium burns in the air, its atoms mix with the gas oxygen to form magnesium oxide. The bright flame produced by flares is produced by this unique combination reaction. 

For examples,

H2 + Cl2 → 2HCl

C + O2 → CO2

Displacement Reaction

An atom or an ion of a compound is replaced by an atom or an ion of another element in this type of reaction. It can be represented as X + YZ → XZ + Y. Displacement reactions are further subdivided into metal displacement reactions and non-metal displacement reactions.

  • Metal Displacement: A metal existing in the compound is displaced by another metal in this type of reaction. These types of reactions are used in metallurgical operations to extract pure metals from their ores. As an example,

CuSO4 + Zn → Cu + ZnSO4

  • Non-Metal Displacement: In this type of reaction, we can detect a hydrogen displacement and, on rare occasions, an oxygen displacement.
  • Single Displacement Reaction: A single displacement reaction, also known as a single replacement reaction, is a type of oxidation-reduction chemical reaction in which one ion or element moves out of a molecule, i.e., one element in a compound is replaced by another.
  • Double Displacement Reaction: Double displacement reactions occur when a portion of two ionic compounds is transferred, resulting in the formation of two new components. This is the pattern of a twofold displacement reaction. Ions precipitate and exchange ions in aqueous solutions, resulting in double displacement processes.

Disproportionation Reactions

Disproportionation reactions are those in which a single reactant is oxidized and reduced. The reaction of hydrogen peroxide, when poured over a wound is one real-life example of such a process. At first glance, this may appear to be a simple breakdown reaction, because hydrogen peroxide decomposes to produce oxygen and water. 

For example,

P4 + 3NaOH + 3H2O → 3NaH2PO2 + PH3

Examples of Redox Reactions

Redox Reactions include both oxidation and reduction reactions. Examples of redox reactions are as follows:

Reaction between Iron and Hydrogen Peroxide

In this redox reaction, H2O2 oxidizes Fe+2 into Fe+3 in the presence of an acid. This results in the formation of Hydroxide ions as shown below:

2Fe22+ + H2O2 + 2H+    →    2Fe3+ + 2H2O

  • Oxidation-half reaction is,

Fe2+    →    Fe3+ + e–

  • Reduction-half reaction is,

H2O2 + 2e–     →     2 OH–

Reaction between Hydrogen and Fluorine

Here in this redox reaction, oxidation occurs at hydrogen and reduction takes place at fluorine. So, Hydrogen and fluorine combine to form Hydrogen Fluoride, as shown below:

H2 + F2     →     2HF

  • Oxidation half of the reaction is, 

H2 → 2H+ + 2e–

  • Reduction-half of the reaction is,

F2 + 2e–     →     2F–
 

Oxidation and Reduction Reaction

Oxidation and Reduction reactions are the basis of Redox reactions so to fully understand the concept of redox reactions, let us learn about oxidation and reduction reactions separately.

What is Oxidation Reaction?

When an atom loses electrons oxidation happens, the oxidation reaction is also defined as the addition of an oxygen atom or the removal of a hydrogen atom. Oxidation number of the atom increases in oxidation reaction.

Some examples of oxidation reactions are:

  • CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l)
  • 2S(s) + O2 (g) → SO2 (g)

What is Reduction Reaction?

When an atom gains electrons reduction happens, the reduction reaction is also defined as the removal of an oxygen atom or the addition of a hydrogen atom. Oxidation number of the atom decreases in reduction reaction.

Some examples of reduction reactions are:

  • 2CH4 (g) + H2 (g) → 2CH6 (g)
  • 2CuCl2 (aq) + H2 (g) → 2CuCl (aq) + 2HCl (aq)

Observing above reactions we can conclude that both oxidation and reduction happens in above reactions. CuCl2 is reduced as electronegative element chlorine is removed from it. While hydrogen is oxidized due to the addition of chlorine, an electronegative element, in the above reaction.

Oxidizing and Reducing Agents

Oxidizing and Reducing agents are defined as,

Oxidizing Agent: Oxidizing agent are substance that gains electrons and is there oxidation number is reduced.

Reducing Agent: Reducing agent are substance that lose electrons and is there oxidation number is increased.

Important Oxidizing Agents

Some Important Oxidizing Agents include 

  • Electronegative elements such as, O3, and X2 (halogens)
  • Oxides of metals and non-metals such as MgO, CuO, CrO3
  • Compounds containing an element in higher oxidized state such as, KMnO4, K2Cr2O7, HNO3

Note: Fluorine is the strongest known oxidizing agent.

Important Reducing Agents

Some Important Reducing Agents include 

  • Metals such as Na, Zn, Fe, Al
  • Hydracids such as HCl, HBr, HI, H2S
  • Metallic hydrides such as NaH, LiH, CaH2
  • Compounds containing elements in lower oxidation state such as FeCl2, FeSO4, SnCl2, Hg2Cl2

Note: Lithium is the strongest reducing agent in the solution, and Cesium is the strongest reducing agent in the absence of water.

Reduction Potential of a Half-Reaction

A redox reaction has two half-reaction and each half reaction has a standard electrode potential. This standard electrode potential equals the voltage produced by the electrochemical cell in which half-reaction is considered as cathode reaction, while standard hydrogen electrode act as anode.

The voltage produced by half-reactions is called their reduction potentials. It is denoted by E°red. For the oxidizing agents stronger than H+ the reduction potential of half reaction is considered to be positive whereas those are weaker than H+ are considered to be negative.

Reduction potentials of some ions include +2.866 V for F2 and -0.763 V for Zn2+.

Identification of Oxidizing and Reducing Agents

Oxidizing and Reducing Agents are identified as,

  • Element which are is in higher oxidation state in a compound behaves as oxidizing agent such as KMnO4, K2Cr2O7, HNO3, H2SO4, HClO4
  • Element which are is in lower oxidation state in a compound behaves as as a reducing agent such as H2S, H2C2O4, FeSO4, SnCl2
  • If high electronegative element is in its highest oxidation state is behaves as oxidizing agent.
  • The compound acts as a reducing agent if a highly electronegative element is in its lowest oxidation state.

Example: Identify the oxidizing agent and reducing agent in the reactions.

Balancing Redox Reaction

There are two ways of balancing the redox reaction. One method is by using the change in oxidation number of oxidizing agent and the reducing agent and the other method is based on dividing the redox reaction into two half reactions-one of reduction and another of oxidation.

Balancing Redox Reaction by Oxidation State Method

This method is based on the fact that the number of electrons gained during the reduction reaction equals the number of electrons lost in oxidation. An example of this is an equation of rusting of iron. The following equation represents rusting of Iron.

4Fe2+ + O2 → 4Fe3+ + 2O2-

Fe2+ is oxidized to Fe3+ Oxidation by gaining 4 electrons (1 by each Fe atom) and two oxygen atoms each losing 2 electrons lead to loss of 4 electrons.

Balancing Redox Reaction by Ion electron Method (Half reaction method)

The steps for balancing redox reaction are mentioned below:

Step 1: Divide the complete reaction into two half-reactions, each called Redox Half Cell. One Redox representing oxidation called Oxidation Half Cell and the other representing reduction is called Reduction Half Cell.

Step 2: First balance elements other than ‘O’ and ‘H’ atoms.

Step 3: In an acidic or neutral medium, balance oxygen atoms by adding H2O molecules and balance H–atoms by adding H+ ions.

OR

In an alkaline medium, the oxygen atom is balanced by adding H2O molecule, and an equal number of ions are added on the opposite side, H+ atoms still unbalanced add OH–

Step 4: Balance the charges by the addition of electrons.

Step 5: Multiply with a suitable integer such that the number of electrons gets cancelled.

Step 6: Add both the half-reactions, similar terms are subtracted, and the final equation is written.

For more details click here, Balancing Redox Reactions

Real-life Applications of Redox Reaction

Redox reactions have numerous industrial and everyday applications. A few of these applications of redox reactions are listed below.

Applications of Redox Reaction in Electrochemistry

The battery used for generating DC current uses a redox reaction to produce electrical energy.

Batteries or electrochemical cells used in our day-to-day life are also based on redox reactions. For example, storage cells are used in vehicles to supply all the electrical needs of the vehicles.

Redox Reaction in Combustion

Combustion is a type of oxidation-reduction reaction, and hence it is a redox reaction. An explosion is a fast form of combustion; hence explosion can be treated as a redox reaction. Even the space shuttle uses redox reactions. The combination of ammonium perchlorate and powdered aluminium inside the rocket boosters gives rise to an oxidation-reduction reaction.

Applications in Photosynthesis

Green plants convert water and carbon dioxide into carbohydrates, defined as photosynthesis. The reaction is given as 6CO2 + 6H2O → C6H12O6 + 6O2

In the above reaction, we can see that carbon dioxide is reduced to carbohydrates while the water gets oxidized to oxygen hence it is a redox reaction. The energy is provided by the sunlight for this reaction. This reaction is a source of food for animals and plants.

Uses of Redox Reaction

Redox Reaction are of great use in our daily life, some examples of redox reaction include:

  • Production of some important chemicals is also based on electrolysis which in turn is based on redox reactions. Many chemicals like caustic soda, chlorine, etc., are produced using redox reactions.
  • Oxidation-Reduction reactions also find their application in sanitizing water and bleaching materials.
  • The surfaces of many metals can be protected from corrosion by connecting them to sacrificial anodes, which undergo corrosion instead. A common example of this technique is the galvanization of steel.
  • The industrial production of cleaning products involves the oxidation process.
  • Nitric acid, a component of many fertilizers, is produced from the oxidation reaction of ammonia.
  • Electroplating is a process that uses redox reactions to apply a thin coating of a material on an object.
  • Electroplating is used in the production of gold-plated jewellery.
  • Many metals are separated from their ores with the help of redox reactions. One such example is the smelting of metal sulfides in the presence of reducing agents.

Related Resources,

Redox Reactions JEE Mains Questions

Q1. The number of electrons involved in the reduction of permagnate to manganese dioxide in acidic medium is

Q2. 5 g of NaOH was dissolved in deionized water to prepare a 450 ml of stock solution. What volume (in ml) of this solution would be required to prepare 500 ml of 0.1 M solution?

Q3. The density of a monobasic strong acid (Molar mass 24.2 g/mol) is 1.21 kg/L. The volume of its solution required for the complete neutralization of 25 ml of 0.24 M NaOH is,—- 10-2 ml (Nearest Integer)

Q4. The volume of 0.2 M aqueous HBr required to neutralize 10.0 ml of 0.001 m aqueous Ba(OH)2 is,

Q5. 2IO3– + xI– + 12H+ –> 6I2 + 6H2. What is the value of x?

Redox Reaction- FAQs

1. What is a Redox Reaction?

The reactions that involve both oxidation and reduction reactions are called redox reactions. 

2. What are Oxidation-Reduction Reactions?

Oxidation-reduction reactions are the chemical reactions that involve transfer of electrons between the reacting species. This led to change in the oxidation state of reactant.

3. What is Redox Reaction Example?

Examples of Redox reactions include,

  • Reaction Between Hydrogen and Fluorine
  • Reaction Between Zinc and Copper Sulphate, and others

4. How to balance Redox Reaction?

Redox reactions are balanced using any of the two methods which include:

  • Oxidation Number method.
  • Ion-Exchange method.

5. What are Oxidizing Agents?

The substance that are readily reduced in a redox reaction are called oxidizing agents. They are electron-accepting species. Oxidation numbers of oxidizing agent decrease in redox reactions. Examples of the oxidizing agent include nitric acid (HNO3) and hydrogen peroxide (H2O2).

6. What are Reducing Agents?

The substance that are readily oxidized in a redox reaction are called reducing agents. They are electron-donating species. Oxidation numbers of reducing agent increases in redox reactions. Examples of reducing agents include lithium(Li) and zinc (Zn).

7. Is every chemical reaction a Redox Reaction?

No, not every chemical reaction is a redox reaction. Non-redox reactions include reactions like double decompositions, acid-base neutralization, and others.

8. What is Redox Reaction of Photosynthesis?

In Photosynthesis, the carbon dioxide molecules obtained from nature is reduced to glucose molecules and the water molecule is oxidised into free oxygen molecule. The redox reaction of Photosynthesis is given as 6CO2 + 6H2O → C6H12O6 + 6O2

9. What we have to study in Redox Reaction Class 10?

In Redox Reaction class 10 we read basic definition of oxidation, reduction, oxidizing agent, reducing agent and application of redox reaction in daily life.

10. What we have to study in Redox Reaction Class 11?

In Redox Reaction Class 11, we study advanced definition of oxidation and reduction, learn to identify the atoms going under oxidation and reduction by calculating charge, learn to balance redox reactions by half cell method and ion electron method.


 

Nernst Equation

Last Updated : 24 Feb, 2022

The electrical potential disparity across the cell membrane of all living cells is called the membrane potential, the inner part of the cell being negative compared to the outside. The magnitude of the membrane potential varies from cell to cell and in an exceptional cell following its functional state. For example, a nerve cell has a membrane potential of -70mv at rest, but the membrane potential drops to about +30mv when excited. The membrane potential at rest is called the resting potential. RMP is basically due to-

  • Uneven distribution of ions across the cell membrane due to its selective permeability.
  • Due to the combined effect of forces acting onions. This origin of RMP is dependent

Selective Permeability of Cell Membrane

The cell membrane is selectively permeable, i.e. it is freely permeable to K+ and Cl–, from the medium to Na+, and is impermeable to proteins and organic phosphates which are negatively charged ions. Usually, the intracellular cation is K+ and major intracellular anions are proteins and organic phosphate. General extracellular cation is Na+ and anion is Cl–. 

The presence of gated protein channels in the cell membrane is responsible for the variable permeability of ions. The forces executive on the ions across the cell membrane Production of variations in the membrane potential. The magnitude of forces acting thenceforward the cell membrane on each ion can be analyzed by the Nernst equation.

Concentration gradient: The Donnan effect results in an uneven distribution of diffuse ions across the cell membrane, in the form of additional diffuse cations, resulting in a concentration gradient.

Electrical gradient: As a consequence of concentration gradient cation K+, will try to disseminate back into ECF from ICF. But it is counteracted by an electrical gradient which will be created due to the impendence of nondiffusible anions within the cell. The membrane potential at which the electric force is equal in magnitude but opposite in the direction of the concentration force is called the equilibrium potential for that ion. The magnitude of the equilibrium potential is determined by the Nernst equation.

Em  = -RT/ZF In Cin/Cout

37 °C at normal body temperature, substituting for the constants (R, T, and F) and converting to the normal logarithm,

Em= -61.5 log Cin/Cout

Nernst equation

The Standard electrode potentials are measured at their standard states when the concentration of the electrolyte solution is fixed as 1M and the temperature is 298 K. Despite this, in actual practice, electrochemical cells do not always have a fixed concentration on the electrolyte solution. The electrode potential depends on the concentration of the electrolyte solution. The Nernst equation gave a relationship between the electrode potential and the concentration of electrolyte solutions, known as the Nernst equation. For a general electrode:

Mn+ + ne–  → M

Ecell = E°cell–(RT/nF)lnQ

Ecell = E°cell–(RT/nF)ln([C]c[D]d / [A]a[B]b)

where,

  • R  = The gas constant  (8.314JK-1mol-1)
  • F = Faraday constant (96,500Cmol-1)
  • T = Temperature in kelvin 
  • Q = Reaction quotient
  • n = Total number of moles of electrons translocate

It should be remembered that when writing the Nernst equation for the overall cell reaction, the log term is the same as the expression for the equilibrium constant for the reaction. The relation of both is similar.

Ecell = E°cell – (2.303 RT/nf) ln([C]c[D]d/ [A]a[B]b)

Ecell = E°cell  + 2.303 RT [A] [B]ln([A]a[B]b/ [C]c[D]d) 

Similarly, for the electrode reaction:

Mn+ + ne–  → M

The Nernst equation is- 

Ecell = E°cell – (2.303 RT/nf)log[1/Mn+]

or Ecell = E°cell  + (2.303RT / nF)log[1/Mn+]

when T = 273K, F=96500 Cmol-1, R=8.314JK-1mol-1 and concentration of solid M is taken as unit

Ecell = E°cell – (0.059/n)log[1/Mn+]

Relationship between equilibrium constant and standard potential of a cell

Ecell = Ecell° – (2.303 RT/nF)log[Kc], [Kc = equilibrium constant]

At equilibrium, Ecell =0

E°cell = (2.303 RT/nF)log[Kc] 

Kc = antilog[nEcell°/0.0591]

Limitations of Nernst Equation

  • The Nernst equation applies exclusively because no current flows through the electrodes. When current flows, the movement of ions at the electrode surface changes, and the conditions of excess potential and resistance loss contribute to the measured potential.
  • At very low concentrations of commutation potential-determining ions, the potential approach ± is found using the Nernst equation. This is corporeally useless, because, underneath such a situation, the commutation current density is reduced and tends to control the electrochemical behaviour of the system more than other effects.
  • Since the active coefficients are close to unity in dilute solutions, the Nernst equation can be expressed in the directly implicit form of the concentration. But in the case of higher concentrations, the actual activities of the ions must be used. This creates complexity for the use of the Nernst equation because estimating the non-ideal activities of ions usually requires experimental measurements.

Sample Questions

Question 1: Will the Eº value change when the coefficients in the chemical equation change?

Answer:

The Eº value does not depend on the coefficient in the chemical equation i.e. when we double or triple the coefficient, the E° value does not change.

For example:

  • Zn²+ 2e– → Zn; E° =-0.76 V
  • 2Zn²+ 4e– → 2Zn; E° =-0.76 V
  • 3Zn²+ 6e– → 3Zn; E° =-0.76 V

In the half-reaction, if the coefficients change, the number of electrons will change to cancel out the effect of the change in n coefficients.

Question 2: Which reference electrode is used to measure the electrode potential of other electrodes?

Answer:

The standard hydrogen electrode is used as a reference electrode whose electrode potential is assumed to be zero. The electrode potential of the other electrode is measured concerning it.

Question 3: Zinc rod is dipped in 0.1M solution of ZnSO4. The salt is 95% dissociated at this dilution at 298 K. Calculate the electrode potential given that E (Zn²+ | Zn) = -0.76 V.

Answer:

The electrode reaction is :

Zn2+ + 2e– ⇆ Zn(s)

According to Nernst equation, at 298 K

E(Zn2+ |Zn)=E° (Zn2+ |Zn)- (0.059 /n) log [ angle n] [Zn]/[Zn2+ (aq)]

E° (Zn2+ |Zn)=-0.76 V, [Zn] = 1,

[Zn2+(aq)]=0.1*95/100=0.095 M

E(Zn2+ |Zn)=-0.76- (0.009/n )log1/-0.095

=-0.76-0.03=-0.79V

Question 4: What advantage do the fuel cells have over primary and secondary batteries?

Answer:

Primary batteries restrain delimited congeries of reactants and are destroyed when the reactants have been consumed. Secondary batteries can be recharged but retract a longer to recharge. The fuel cell is conducted consecutive as long as reactants are recoupment to it and products are continuously removed.

Question 5: How will the pH of the brine (aq. NaCl solution) be affected when it is electrolyzed?

Answer:

Since NaOH is formed during electrolysis, the pH of the brine solution will increase


Conductance of Electrolytic Solutions

Last Updated : 25 Sep, 2022

Electrochemistry is a branch of chemistry, and it deals with the study of the production of electricity from the energy released during spontaneous chemical reactions and the use of electrical energy to bring about non-spontaneous (requiring the input of external energy) chemical transformation. Electrochemistry deals with how much chemical energy produced in a redox reaction can be converted into electrical energy. A redox reaction is that in which oxidation and reduction take place simultaneously. The arrangements used to bring about the chemical transformations are called electrochemical cells. The cells are used to convert chemical energy into electrical energy and electrical energy into chemical energy.

Types of Electrochemical Cells

There are two types of electrochemical cells:

  • Electrolytic cells: Electrolytic cells are used to convert electrical energy into chemical energy. 
  • Galvanic or voltaic cells: Galvanic or voltaic cells are used to convert chemical energy into electrical energy. These galvanic or voltaic cells are also called electrochemical cells. Galvanic cells are further classified into two types they are Chemical cells and Concentration cells. 

Chemical cells

Chemical cells are those in which electrical energy is produced only due to chemical changes occurring within the cell, and no transfer of matter takes place. Example: Batteries. 

Concentration cells

Concentration cells in which electrical energy is produced due to physical changes involving the transfer of matter from one part of the cell to the other. For example – Standard hydrogen electrode

The three main aspects of study in electrochemistry are – Electrolysis or Electrolytic cells, Galvanic or voltaic or Electrochemical cells, and Electrolytic conduction.

Importance of Electrochemistry

Electrochemistry has great importance in everyday life. It has great theoretical and practical importance. There are many examples that indicate the great importance of electrochemistry-

  • Many numbers of metals and chemicals are commercially produced by electrochemical methods. Metals like Na, Ca, Mg, etc., and chemicals like NaOH, Chlorine, Fluorine, etc.
  • Batteries and cells are used in various gadgets or instruments. Example – torches, calculators, remotes, etc. These are used to convert chemical energy into electrical energy.
  • The sensory signals are sent to the brain through the cells, and communication is also possible via these electrochemical processes.
  • Electrochemically, reactions carried out are generally energy efficient and less polluting.
  • Energy storage, energy conversion, sensing, etc., have a great role in electrochemistry.
  • The coating of objects with metals or metal oxides through electrodeposition has a role in electrochemistry.
  • Organic electrosynthesis and industrial electrolysis also have a great role in electrochemistry. These processes are feasible because of electrochemistry.

Conductance of Electrolytic Solutions 

Conductors are those substances that allow electricity to pass through them, whereas substances that do not allow electricity to pass through are called insulators. Conductors are divided into two classes,

  • Electronic conductors: These are those which conduct electricity without undergoing any decomposition. These are called electronic conductors. The conduction, in this case, is due to the flow of electrons. Example – Metals, Graphite, and certain minerals.
  • Electrolytic conductors: These are those which undergo decomposition when current passes through them. These are called Electrolytic conductors. Example: Solution of acids, bases, salts, and salts in water. In this case, the flow of electricity is due to the movements of ions. Hence, electrolytic conductance is also called ionic conductance.

Metallic conductors

Metallic conductors are those in which the flow of electricity is due to the flow of electrons, i.e., there is no flow of matter. In metallic conductors, the flow of electricity takes place without the decomposition of the substances. The conductance depends on the structure and density of metal as well as the number of valance electrons per atom. The electrical conduction decreases with the increase in temperature because Kernels start vibrating, which produces a hindrance in the flow of electrons. The resistance offered by metal is due to vibrating kernels.

Electrolytic Conductor  

In electrolytic conductors, the flow of electricity is due to the movement of ions, and hence there is no flow of matter. The flow of electricity takes place, accompanied by the decomposition of the substance. The electrical conduction increase with the increase in temperature. This is generally due to an increase in dissociation or a decrease in interionic attraction. The resistance shown by electrolytic solution is due to factors like interionic attractions, the viscosity of solvents, etc.

Electrical resistance and conductance 

Resistance (R) is the obstruction to the flow of electric current through the conductor. It is directly proportional to its length and inversely proportional to the area of cross-section (A). And according to ohm’s law, two ends of a conductor are applied with voltage (E), and current (I) flows through it. Then, the resistance of a conductor is:

R = E/I   

Or       

R = ρ l/A 

Resistance (R) is the obstruction to the (R = E/I) flow of electric current through the conductor. Resistance is directly proportional to its length ( l ) and inversely proportional to its area of cross-section (A). The constant of proportionality ρ (rho) is called specific resistance or resistivity. Resistance is measured in ohm, which in terms of SI base units is equal to (kgm2/S3A2). The S.I. unit of resistivity is ohm meter.

Conductance (G) is the inverse of resistance or the reciprocal of resistance is called conductance, and it is denoted by G. The unit of conductance is ohm inverse or reciprocal ohm or siemens or mhos.   

G = 1/R   

Or   

G = A/ρl = κ A/l

The S.I. unit of conductance is siemens, represented by the symbol S, and is equal to ohm-1.

Specific, Equivalent, and Molar conductivities

Specific Conductivity: It is the conductance (G) of a one-centimeter cubic solution of the electrolyte. It is denoted by (κ), i.e., kappa. It is the conductance of the solution of one-centimeter length (l) and having one square meter area (a) of a cross-section. And specific conductivity is the reciprocal of resistivity is known as conductivity. Its unit is Sm-1 or ohm-1cm-1.

κ = [l × G) / a]

Equivalent Conductivity: Equivalent conductivity of a solution at a dilution (v) is defined as the conductance of all the ions produced from one gram equivalent of electrolyte dissolved in v centimeter cubic of the solution when the distance between the electrodes is one centimeter, and the area of electrodes is so large that whole of the solution is contained between them. It is represented by lambda of equivalent conductivity. Unit of equivalent conductivity is siemens meter square per equivalent.  And  Equivalent conductivity = Specific Conductivity × V; V is the volume of solution containing one gram equivalent of the electrolyte is V centimeter cubic. If the solution has a concentration of c gram equivalent per liter, then the volume of the solution containing one gram equivalent will be 1000/c i.e., V = 1000/c

Λeq   = [κ × (1000/ceq )] = [κ × (1000/Normality)]

Its unit is Sm2eq-1 or Scm2eq-1

Electrolysis

 

Molar Conductivity: Molar Conductivity of a solution at a dilution (v) is defined as the conductance of all ions produced from one mole of the electrolyte dissolved in the v centimeter cubic of the solution when the electrodes are one centimeter, and the area of the electrode is so large that whole solution is contained between them. It is also represented by lambda of molar conductivity. Its unit is siemens meter square per mole.                                        

Molar conductivity = Specific Conductivity × V

Conductivity Cell (G*) is used to measure the resistance of an ionic solution. It consists of two platinum electrodes coated with platinum black. These have an area of cross-section equal to A and are separated by distance (l).  

Cell constant (G*) = l / A ; 

Conductivity (K) = Conductance (G)  × Cell constant ( G*)

Its unit is Sm2mol-1. 

Strong Electrolyte: Those electrolytes which dissociate almost completely in the aqueous solution or in a molten state are called strong electrolytes. Example – HCl, Sulfuric acid, nitric acid, etc.

Weak Electrolyte: Those electrolytes which have a low degree of dissociation and hence conduct electricity to a small extent are called weak Electrolytes. Example – Ammonium hydroxide, Calcium hydroxide, etc. 

Kohlrausch Law of Independent migration of ions: Limiting molar conductivity of an electrolyte is the sum of the individual contribution of the anion and cation of the electrolyte. The limiting molar conductivity of an electrolyte is the sum of limiting ion conductivities of the cation and the anion, each multiplied by a number of ions present in the one formulae unit of electrolyte.

Degree of dissociation of weak electrolyte: It is represented by alpha. It is defined as the Molar conductivity of a solution at any concentration (c) divided by limiting molar conductivity.

Dissociation Constant of a Weak electrolyte: The dissociation constant of a weak electrolyte is directly proportional to the concentration of the solution and the square of the degree of dissociation of the weak electrolyte.

Solved Examples on Conductance of Electrolytic Solutions

Example 1: The resistance of a conductivity cell containing 0.001M KCl solution at 298K is 1500 ohm in a conductivity cell. If the cell constant of the cell is 0.367 per cm, calculate the molar conductivity of the solution.

Solution:

Cell constant = Conductivity/Conductance = Conductivity × Resistance

= (0.146 × 10-3) Scm-1 × 1500 ohm 

= 0.219cm-1

Example 2: The Conductivity of 0.20 m solutions of KCl at 298K is 0.0248 Scm-1. Calculate its molar conductivity.

Solution: 

Molar conductivity = (κ × 1000)/Molarity = [(0.0248Scm-1 × 1000cm3 L-1)/ 0.20 molL-1]

= 124 Scm2 mol-1

Example 3: The electrical resistance of a column of 0.05mol/L NaOH solution of diameter 1 cm and length 50 cm is 5.55 × 103 ohm. Calculate resistivity, conductivity, molar conductivity.

Solution:

Area = πr2 = 3.14 × 0.52 cm = 0.785 cm2 = 0.785 ×10-4 m2, ρ = 5.55 × 103 ohm 

R = ρ l/A = [(5.55× 103 ohm × 0.785cm2)/50 cm] 

= 87.135 ohm cm

Conductivity = κ = 1/ρ = (1/87.135)Scm-1  = 0.01148 Scm-1 

Molar Conductivity = [(κ × 1000)/c] cm3L-1 = (0.01148 Scm-1 × 1000 cm3L-1)/0.05molL-1

= 229.6 Scm2 mol-1

FAQs on Conductance of Electrolytic Solutions

Question 1: How does the conductivity of the solution vary with concentration?

Answer: 

Conductivity is conductance between two opposite faces of the one-centimeter cube. On dilution, the number of ions per centimeter cubic decreases; therefore, conductivity decrease on dilution. 

Question 2: How does the Molar Conductivity of Strong and Weak electrolytes vary with concentration?

Answer:

In the case of Strong Electrolytes, the molar conductivity increases slightly with dilution as the mobility of ions increases. In case of weal electrolytes the degree of ionization increases with dilution. Therefore, there is a large increase in molar conductivity with dilution.

Question 3: What is the application of Kohlrausch law?

Answer: 

Application of Kohlrausch law is,

  • Calculation of molar conductivity at infinite dilution for weak electrolytes: The molar conductivity of a weak electrolyte at infinite dilution cannot be determined experimentally. Firstly, because the conductance of such solution is low, and secondly, because the dissociation of such electrolyte is not complete even at very high dilution. 
  •  Calculation of Degree of Dissociation 
  • Calculation of Dissociation constant of a weak electrolyte 
  • Calculation of solubility of a sparingly soluble salt – Salts such as AgCl, Barium sulfate, Lead sulfate, etc., which dissolved to a very small extent in water called sparingly  soluble salt.
  • Calculation of ionic product of water

Question 4: Explain Metallic Conductor.

Answer:

In a metallic conductor, the flow of electricity takes place without the decomposition of the substance. Here, the flow of electricity is due to the flow of electrons only. The electrical conduction decrease with increase of temperature. This is because Kernels start vibrating, which produces a hindrance in the flow of electrons.

Question 5: Explain Electrolytic Conductor.

Answer:

In an Electrolytic Conductor, the flow of electricity takes place accompanied by the decomposition of the substance. Here, flow of electricity is due to the movement of ions. The electrical conduction increase with increase of temperature. This is generally due to an increase in dissociation or decrease in dissociation or a decrease in the interionic attraction.



 

Variation of Conductivity and Molar conductivity with Concentration

Last Updated : 01 May, 2022

Electrochemistry is the study of chemical reactions that occur in a solution at the interface of an electron conductor (the electrode: a metal or a semiconductor) and an ionic conductor (the electrolyte). Electron transfer occurs between the electrode and the electrolyte or species in solution in these reactions.

Conductivity

A solution’s conductivity is defined as the conductance of a solution with a length of 1 cm and a cross-sectional area of 1 sq. cm. Conductivity, or particular conductance, is the inverse of resistivity. The letter k is used to signify it. If p stands for resistivity, we can write:

K= 1/p 

The conductivity of a solution at any given concentration is equal to the conductance (G) of one unit volume of solution held between two platinum electrodes with the same cross-sectional area and separated by the same distance.

i.e., 

G = K × a/l = K × l = K

(Since a = 1, l = 1)

For both weak and strong electrolytes, conductivity diminishes as concentration decreases. This is because as the concentration of a solution declines, the number of ions per unit volume that carry the current in the solution reduces.

Molar Conductivity

The conductance of volume V of a solution containing 1 mole of the electrolyte kept between two electrodes with the area of cross-section A and distance of unit length is the molar conductivity of a solution at a particular concentration.

Am = K × A/l

Now, l = 1 and A = V (volume containing 1 mole of the electrolyte).

Am = KV

With a decrease in concentration, molar conductivity rises. This is due to the fact that when a solution containing one mole of electrolyte is diluted, the total volume V of the solution increases.

The following graph depicts the fluctuation of Am with √c for strong and weak electrolytes:

 

Variation of Molar Conductivity with Concentration for Strong and Weak Electrolytes

  • Variation of Molar Conductivity with Concentration for Strong Electrolytes

Molar conductivity grows slowly with dilution in strong electrolytes, and it has a tendency to approach a limiting value as the concentration approaches 0, i.e. when the dilution is infinite. Molar conductivity at infinite dilution is the molar conductivity as the concentration approaches 0 (infinite dilution).  It is denoted by Am°

Am = Am°, when C ⇢ 0 (at infinite dilution)

The expression for the change of molar conductivity with concentration might be used.

Am = Am° − AC1/2

where 

  • A is constant and A° stands for molar conductivity at infinite dilution. 

This equation, known as the Debye Huckel Onsager equation, holds true at low concentrations.

  • Variation of Molar Conductivity with Concentration for Weak Electrolytes

When opposed to strong electrolytes, weak electrolytes dissociate to a far smaller level. As a result, when compared to strong electrolytes, the molar conductivity is low.

However, the variation of Am with C1/2 is so enormous that the extrapolation of Am against C1/2 plots cannot yield molar conductance at infinite dilution ( Am°).

Variation of Molar Conductivity with Concentration

(A) Conductance behavior of weak electrolytes

The degree of dissociation with dilution determines the number of ions supplied by an electrolyte in a solution. The degree of dissociation rises as dilution increases, and molar conductance increases as a result. The limiting value of molar conductance (Am) corresponds to a degree of dissociation of 1, which means that the electrolyte completely dissociates.

At every concentration, the degree of dissociation may therefore be estimated.

α = Amc / Am°

where α represents the degree of dissociation, Amc represents the molar conductance at concentration C, and Am° represents the molar conductance at infinite dilution.

  • Conductance behavior of strong electrolytes

For strong electrolytes, there is no increase in the number of ions with dilution because strong electrolytes are completely ionized in solution at all concentrations.

Interionic forces are strong forces of attraction between ions of opposing charges in concentrated solutions of strong electrolytes. In concentrated solutions, the ions’ conducting capacity is reduced due to these interionic interactions. The ions grow farther apart as a result of dilution, and interionic forces diminish. As a result, molar conductivity rises as the solution is diluted. When the solution’s concentration is exceedingly low, interionic attractions are minimal, and molar conductance approaches the limiting value known as molar conductance at infinite dilution. This number is unique to each electrolyte.

Sample Questions 

Question 1: What effect does a solution’s concentration have on its specific conductivity?

Answer:

The specific conductivity decreases as the concentration decreases. This is because the number of energized ions per unit volume  in a solution decreases with dilution. Therefore, concentration and conductivity are directly proportional to each other.

Question 2: Explain why the Cu+ ion is not stable in aqueous solutions?

Answer:

Cu2+ is more stable in aqueous media than Cu+. This is because, while removing one electron from Cu+ to Cu2+ requires energy, the high hydration energy of Cu2+ compensates for it. As a result, the Cu+ ion is unstable in an aqueous solution. Cu2+ and Cu are disproportionately produced.

2Cu+(aq) ⇢ Cu2+(aq) + Cu(s)

Question 3:  The molar conductivity of 0.025 mol L-1 methanoic acid is 46.1 S cm2 mol-1. Calculate its degree of dissociation and dissociation constant. Given λ0(H+) = 349.6 S cm2 mol-1 and  λ0(HCOO–) = 54.6 S cm2 mol.

Answer:

Given that,

C = 0.025 mol L-1

Am = 46.1 S cm2 mol-1

 λ0(H+) = 349.6 S cm2 mol-1

 λ0(HCOO−) = 54.6 S cm2 mol-1

Am°(HCOOH) =  λ0(H+) +  λ0(HCOO−) 

= 349.6 + 54.6

= 404.2 S cm2 mol-1

Now, degree of dissociation:

α = Am(HCOOH)/Am°(HCOOH)

= 46.1/404.2

= 0.114(approx.)

Thus, dissociation constant:

K = c×α2/(1−α)

= (0.025 mol L-1)(0.114)2/(1 − 0.114)

= 3.67×10-4 mol L-1

Question 4:  The conductivity of 0.20M solution of KCl at 298K is 0.0248 S/cm. Calculate its molar conductivity 

Answer:

Given:-

 K= 0.0248 S/cm

C= 0.20M

Am = K×1000/C

= (0.02481000)/0.2

= 124 S cm2 mol-1

Question 5: Prove that “Molar conductivity increases with a decrease in concentration”.

Answer:

Keeping length is equal to one.  

Multiply length ‘l’ in numerator and denominator

Am = K×A/l

Keeping length is equal to one.  

Multiply length ‘l’ in numerator and denominator

Am = K×A/l × l/l
As l=1 and A×l =V volume 

Am = KV

The concentration is low, the volume grows, and the cross-sectional area expands. As a result, both the volume and the molar conductivity rise. When the concentration is low, the volume is raised.

Because ‘K’ is related to conductivity, when concentration rises, ‘K’ rises while ‘V’ falls. When the concentration drops, ‘K’ drops as well, while ‘V’ rises. The change in the value of ‘V’ is significantly bigger than the change in the value of ‘K.’




 

Equivalent Conductance Formula

Last Updated : 19 Dec, 2023

Electrochemistry includes the concept of equivalent conductance, which is the conductance of a volume of solution containing one equivalent of an electrolyte. Let’s study the idea of the equivalent conductance formula.

Equivalent Conductance

The term “equivalent conductance” refers to the conductance (or “conducting power”) of all the ions (of a solution) created by dissolving one gram equivalent of an electrolyte in a specific solution.

We can state that an electrolytic solution’s conductance is influenced by the ion concentration present in the solution. Having comparable outcomes for various electrolytes is beneficial. It’s represented by the symbol ∧e. From specific conductance, equivalent conductance is computed.

Unit of Equivalent Conductance is ohm-1cm2eq-1.

Equivalent Conductance Formula is as follows:

∧e = K × V

Where,

  • ∧e = Equivalent Conductance,
  • K = Specific Conductance (Reciprocal of Specific Resistance),
  • V = Volume (in ml) of 1 gm-equivalent electrolyte.

Also,

∧e = (K × 1000) / N

Where,

N = Normality

Derivation of Equivalent Conductance

Conductance of V cm3 = ∧e

Conductance of 1 cm3 = K

Therefore, ∧e = K × V  …(Equation 1)

We are aware that the equation below provides information about a solution’s normality (N).

N = (n/V) × 1000

∴ V = (1000 × n) / N

Number of equivalents, n = 1, for the electrolytic solution mentioned above.

V = 1000 / N

Substitute value of V in Equation 1,

∴ ∧e = (K × 1000) / N

Conductors and Insulators

Conductors are substances that make it simple and unhindered for electrons to go from one end to the other. Conductors contain electric charges in the form of electrons, which facilitates the electrons’ free movement.

On the other hand, insulators are the kind of substances that obstruct the easy movement of electrons from one end to other. Any charge that is transferred through an insulator only settles at the point where the two materials first come together; it does not expand outward.

Equivalent conductance at infinite dilution

The value of equivalent conductance rises as the solution becomes more diluted as ionization, or the amount of ions in a solution, rises. However, there comes a point where further dilution of the solutions is impossible, meaning it has no impact whatsoever on the solution’s concentration. The infinite dilution is the name given to the entire idea when diluting ceases.

Since a solution already has the maximum amount of solvent that may be added, infinite dilution is the condition in which no further concentration can be achieved with any amount of dilution. In this infinite dilution condition, all ions are fully dissociated.

Kohlrausch’s Law

According to Kohlrausch’s law, each ion contributes significantly to the equivalent conductance of the electrolyte at infinite dilution when dissociation is complete, regardless of the type of ion it is associated with. The value of equivalent conductance at infinite dilution for any electrolyte is the sum of the contributions of its constituent ions (cations and anions). We can therefore interpret it to mean that “conductivity of an electrolyte’s ions at infinite dilution is constant and does not depend on the nature of co-ions.”

λ∞eq = λ∞c + λ∞a

Where,

  • λ∞eq = Molar conductivity at infinity of dilution,
  • λ∞c = Conductivity of cation at infinity of dilution,
  • λ∞a = Conductivity of anion at infinity of dilution.

The term “limiting molar conductivity” refers to the molar conductivity that exists when the electrolyte concentration is almost zero.

When held between two electrodes with a unit area of cross-section and a unit distance, the volume of the solution that conducts and also contains one mole of electrolyte is known as the molar conductivity. Molar conductivity rises as concentration falls. The volume that makes up one mole of electrolytes increases, which results in an increase in the molar conductivity. When the electrolyte concentration becomes close to zero, the molar conductivity is known as the limiting molar conductivity.

Uses of Kohlrausch’s Law

  1. The molar conductivity at infinite dilution for the weak electrolytes is determined using Kohlrausch’s law. Calculating the molar conductivity of weak electrolytes at infinite dilution is highly challenging or impossible. due to the extremely low conductance of these sorts of solutions and the fact that the dissociation of these electrolytes is incomplete even at high dilutions.
  2. The solubility of a moderately soluble salt is determined using Kohlrausch’s law. Some salts are referred to as weakly or sparingly soluble salts because they only slightly dissolve in water. For instance, silver chloride, lead sulfate, barium sulfate, etc.
  3. The term “limiting molar conductivity” refers to the molar conductivity that exists when the electrolyte concentration is almost zero. We can calculate the limiting molar conductivity of an electrolyte using Kohlrausch’s law.

Molar Conductivity

Molar conductivity, which can be determined by a solution’s ionic strength or salt concentration, is the conductance of a solution containing one mole of electrolyte. It is therefore not a constant. 

Molar conductivity, then, is the sum of the conductivities of all the ions produced when a mole of an electrolyte is dissolved in a solution. The ability of an electrolyte to transmit electricity in a solution is generally assessed using the property of an electrolyte solution called molar conductivity. It is therefore not a constant. The unit of Molar conductivity is Sm2mol-1.

The following expression is used to numerically represent molar conductivity:

μ = K/C

Where,

  • K = Specific Conductivity,
  • C = the concentration of moles per liter.

Factors Affecting Equivalent Conductivity

  • Temperature: As temperature rises, more ions are produced, which results in an increase in an electrolyte solution’s conductance.
  • Strong electrolytes completely ionize, producing more ions as a result of which they have higher conductivities.
  • On the other hand, weak electrolytes only experience partial ionization, which results in low conductivity in their solutions.
  • Ionic size and mobility: As an ion’s size grows, so does its mobility, and its conductivity also decreases.
  • Because of the solvent’s composition and viscosity, ionic mobility is reduced in more viscous solvents. The conductivity consequently declines.

Sample Questions

Question 1: What Does Equivalent Conductance Mean in Chemistry?

Answer:

Equivalent conductance of an electrolyte is defined as the conductance of a volume of solution containing one equivalent weight of dissolved substance when placed between two parallel electrodes spaced 1 cm apart and big enough to hold the entire solution between them.

Question 2: Write Two Factors Affecting Equivalent Conductivity.

Answer:

Factors Affecting Equivalent Conductivity:

  1. Strong electrolytes fully ionize, resulting in the production of more ions and higher conductivities.
  2. Ionic size and mobility: An ion’s mobility increases with its size while its conductivity decreases.

Question 3: List the situations in which the independent ion migration law of Kohlrausch is applicable.

Answer:

For weak electrolytes, it is used to determine the limiting molar conductivity, level of dissociation, and dissociation constant. It is also employed in the computation of the salt’s solubility.

Sample Problems

Problems 1: A 0.7 N salt solution put between two platinum electrodes separated by 2 cm and covering an area of 6 cm2 has a 25-ohm resistance. calculate equivalent conductivity.

Solution:

Since,

∧e = (K × 1000) / N

∴ ∧e = 1/25 × 2/6 × 1000/0.7

∴ ∧e = 0.04 × 0.33 × 1428.5

∴ ∧e = 18.8562 ohm-1cm2eq-1

Problems 2: A salt solution in N/10 is found to have a resistance of 1.2 × 103 ohms. Calculate the solution’s equivalent conductance. 1.5 cm-1 is the cell constant.

Solution:

Since,

∧e = (K × 1000) / N

K = cell constant × conductance

∴ K = 1.5 × (1/1.2 × 103)

∴ K = 1.25 × 10-3

∧e = K × 1000 / N

∴ ∧e = 1.25 × 10-3 / (1/10)

∴ ∧e = 1.25 × 10-3 / 0.1

∴ ∧e = 12.5 × 10-3 ohm-1cm2eq-1

Problems 3: Calculate the volume of the solution if the equivalent conductance is 10.255 ohm-1cm2eq-1 and the specific conductance is 2.17 cm-1ohms-1.

Solution:

Since,

∧e = K × V

∴ V = ∧e / K

∴ V = 10.255 / 2.17

∴ V = 4.7258 cm3

Problems 4: Calculate the Equivalent conductance if the specific conductance is 1.83 cm-1ohms-1 and the volume of the solution is 3.91 ml.

Solution:

Since,

∧e = K × V

∴ ∧e = 1.83 × 3.91

∴ ∧e = 7.1553 ohm-1cm2eq-1

Problems 5: It is discovered that a salt solution in N/100 has a resistance of 5.21 × 103 ohms. Calculate the equivalent conductance of the solution. The cell constant is 3.71 cm-1.

Solution:

Since,

∧e = (K × 1000) / N

K = cell constant × conductance

∴ K = 3.71 × (1/5.21 × 103)

∴ K = 0.7120 × 10-3

∧e = K × 1000 / N

∴ ∧e = 0.7120 × 10-3 / (1/100)

∴ ∧e = 0.7120 × 10-3 / 0.01

∴ ∧e = 71.2 × 10-3 ohm-1cm2eq-1

Problems 6: Calculate the Equivalent conductance if the specific conductance is 3.11 cm-1ohms-1 and the volume of the solution is 2.95 ml.

Solution:

Since,

∧e = K × V

∴ ∧e = 3.11 × 2.95

∴ ∧e = 9.1745 ohm-1cm2eq-1

Problems 7: If the equivalent conductance is 2.916 × 103 ohm-1cm2eq-1 and the specific conductance is 2.87 cm-1ohms-1 then calculate the volume of the solution.

Solution:

Since,

∧e = K × V

∴ V = ∧e / K

∴ V = 2.916 × 103 / 2.87

∴ V = 1.0160 × 103 cm3

Problems 8: Calculate λm for Cacl2 if λCa2+ = 137.0 Scm2mol-1 and λcl– = 65.9 Scm2mol-1.

Solution:

Since,

λom Cacl2 = λo Ca2+ + 2λocl–

∴ λom Cacl2 = 137 + (2 × 65.9)

∴ λom Cacl2 = 137 + 131.8

∴ λom Cacl2 = 268.8 Scm2mol-1

Problems 9: Calculate λom for MgSO4 if λoMg2+ = 119 Scm2mol-1 and λoSO42- = 241 Scm2mol-1.

Solution:

Since,

λom MgSO4 = λo Mg2+ + λo SO42-

∴ λom MgSO4 = 119 + 241

∴ λom MgSO4 = 360 Scm2mol-1


 

Electrolysis

Last Updated : 19 Dec, 2023

Electrolysis is the process of decomposing the ionic compound into its constituent elements by passing the electric current into the solution of the ionic compound. The concept of electrolysis was first given by the famous scientist of the 19th century Michael Faraday. It is a chemical process that uses electrical energy to bring changes in the chemical reaction. Electrolysis is used to separate components of the ionic compounds.

In this article, we will learn about, electrolysis, its process, faraday’s law of electrolysis and others in detail in this article.

Electrolysis Definition

The process of decomposing the ionic compound into its constituent elements using electric current is called Electrolysis. In electrolysis, the ionic compound is dissolved into the aqueous medium and then electricity is passed through the solution using the electrodes then the ionic compound breaks into its constituent elements and cations are collected at the cathode and anions are collected at the anode.

The cell which converts electrical energy into chemical energy is called an electrolytic cell. 

This redox reaction occurs at electrodes specifically mentioning oxidation occurs at the anode and is a positive plate while reduction occurs at the cathode and is a negative plate. In electrolytic cells, electrical energy is used to perform non-spontaneous chemical reactions and the process that takes place in an electrolytic cell is called Electrolysis.

Electrolysis

 

Electrolytic Process

In the Electrolysis Process, the exchange of ions takes place by the electric current passing through the circuit. By allowing current to pass through the solution we force cations to get attached to the cathode of the electrolytic cell and anions to get attached to the anode of the electrolytic cell. 

We can understand the electrolytic process with the help of NaCl example, 

Electrolysis of NaCl Solution

The aqueous solution of NaCl has Na+, Cl–, H+, and OH– ions present in it. Now when the electrodes are introduced to this solution and the electricity is passed through it that allows Na+, and H+, ions to move to the negatively charged electrode, (cathode) and the Cl–, and OH–  ions move to the positively charged electrode, (anode). 

NaCl ⇌ Na+ + Cl–

H2O ⇌ H+ + OH–

At Cathode: 

  • Na+ + e– → Na

At Anode: 

  • 2Cl– →  Cl2 + e–  

Now as a result of this electrolysis solution the Sodium metal is deposited on the cathode of the electrolytic cell and the chlorine is released at the anode.

Cell Potential

The least potential required by the electrolytic cell to complete the process of electrolysis depends on the mobility of the ions in the aqueous solution. If the ions have very low mobility then the cell potential of the electrolytic cell must be high to facilitate the movement of the ions in the aqueous solution. 

Whereas if the mobility of the ions is very high in their aqueous solution then an electrolytic cell with low electric potential is used as it can easily facilitate the movements of ions in the electrolysis process.

The cell potential of an electrolytic cell is the sum of the standard oxidation potential and the standard reduction potential of the cell.

Faraday Laws of Electrolysis

The Faraday laws of electrolysis are the basic laws of electrolysis that provide information about the mass of the substance and the charge in the electrolysis process. There are two Faraday Laws of Electrolysis that are,

  • Faraday’s First Law of Electrolysis
  • Faraday’s Second Law of Electrolysis

Now let’s learn about the Faraday Laws of Electrolysis in detail.

Faraday’s First Law of Electrolysis

Faraday’s First Law of Electrolysis states that “the mass of the substance that undergoes electric current is directly proportional to the charge supplied.”

We know that,

i = Q/t, then Faraday’s First Law of Electrolysis is represented as,

m ∝ Q 

m = ZQ 

m = Zit

where,

  • m is the mass of the substance that undergoes electrolytic process
  • Q is the charge associated with the electrolytic cell
  • i is the current measured in Ampere (A)
  • t is the time measured in sec
  • Z is the constant of proportionality called Electrochemical Equivalent.

Faraday’s Second Law of Electrolysis

According to the second law of electrolysis, the amount of electrolyte deposited at the electrodes is the directly proportional equivalent weight of the material, if the amount of electricity passing through the solution is constant.

 W1/E1 = W2/E2 = W3/ E3 …….

where

  • W1 is the mass of first substance, W2 is the mass of the second substance and so on…
  • E1 is the Equivalent Weight of the first substance, E2 is the Equivalent Weight of the second substance and so on…

Product of Electrolysis

The oxidising and the reducing species present in the electrolytic cell are used to find the product of the electrolysis. If in an aqueous solution, we have more than one cation or anion then each ion will not be reduced or oxidised. The reaction with more redox potential will be reduced or oxidised in comparison to others.

For example, the electrolysis of aqueous sodium chloride gives different products such as,

  • Hydrogen and Chlorine
  • Hydrogen and Oxygen
  • Hydrogen, Oxygen and Chlorine

Electrolysis of Sodium Chloride (NaCl) solution also called Brine is discussed below,

  • At Anode: Chlorine Gas (Cl₂) is produced
  • At Cathode: Hydrogen Gas (H₂) is produced

In the solution, Sodium Hydroxide (NaOH) is formed.

Factors Affecting Electrolysis

Various factors affecting electrolysis are discussed below,

  • Nature of the Electrolyte
  • Nature of the Electrode
  • Voltage at the Electrodes

Now let’s learn about them in brief.

Nature of the Electrolyte

In the electrolysis process, the charged particles called ions dissolved in the aqueous solution move under the influence of the electric potential, and so the nature of the electrolyte plays important in the electrolysis process. If we take an electrolyte that dissolves easily in the aqueous solution then it facilitates the process of electrolysis.

Nature of the Electrode

Nature of the electrode plays an important role in the electrolysis process taking different electrodes in the same electrolyte solution gives different outputs.

The aqueous solution of copper sulphate solution on electrolysis gives the following results,

At cathode: (Reduction)

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

At anode: (Oxidation)

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

Here, we see that if copper is taken at the cathode we get a different reaction and if copper is taken at the anode we get a different reaction.

Voltage at Electrodes

The redox potential of the electrolyte plays an important role in the electrolysis process. If the redox potential of the electrolysis reaction is more than the thermodynamic potential of the electrolysis reaction the reaction becomes unfavourable and the product of electrolysis changes.

If we perform the hydrolysis of aqueous sodium chloride, at the anode there are two oxidation reactions possible. The reduction potential of water and chloride is +0.82V and -1.36V, respectively in both these reactions and they are represented as,

  • 2H2O→O2(g) + 4H+ + 4e–    (E° = -0.82 V)
  • 2Cl– → Cl2 + 2e–                 (E = -1.36V)

Here we observe that the oxidation of water is more positive and hence is more feasible, so oxygen is evolved at the anode. But the evolution of the oxygen resulted in an overvoltage of -0.6 V making the voltage of oxidation of water to be -1.42

Applications of Electrolysis

Electrolysis is one of the most important processes used in electrochemistry. It is used for various purposes and some of its applications are,

  • Caustic soda is prefabricated by electrolysis of sodium chloride solution
  • Manufacture of O2 and H2 by Electrolysis of Water
  • Determination of Equivalent Weight of Substances
  • Purification of Metals
  • Electroplating for Corrosion Resistance, Ornaments etc.

Learn more about Refining of Metals

What is Electrolytic Cell?

Electrolysis is the procedure of decomposition of an electrolyte by the passage of electricity throughout its aqueous solution or the molten state of an electric current. This cell is utilized to perform electrolysis which is electrolyte cells.

Water, for example, can be electrolyzed (with the help of an electrolytic cell) to produce gaseous oxygen and hydrogen. This is accomplished by utilising the flow of electrons (into the reaction environment) to overcome the non-spontaneous redox reaction’s activation energy barrier.

The following are the three major components of electrolytic cells:

  • Cathode: It is negatively charged for electrolytic cells
  • Anode: It is a type of electrode that is (which is positively charged for electrolytic cells)
  • Electrolyte: The electrolyte serves as a conduit for electrons to flow between the cathode and the anode. Water (containing dissolved ions) and molten sodium chloride are common electrolytes in electrolytic cells.

Read More,

Solved Examples on Electrolysis

Example 1: How many coulombs are needed for 40.5 g of aluminium to react when the electrode is:

Al3+ + 3e– ⇒ Al

Solution:

1 mol of Al requires 3 mol of electrons or 3 × 96500 C

1 mol of Al = 27g

27g of Al require =3 × 96500 C

40.5g of Al require =(3*96500C × 40.5)/27 = 434,250 C

Example 2: In the electrolysis of acidic water, it is desired to obtain hydrogen at 1cc sec at the STP position. What should be the current pass?

Solution:

2H+ + 2e– ⇒ H2

1 mol of H2 or 22400 cc of H2 at STP requires = 2 × 96500 C

1cc of H2 at NTP requires = (2×96500)/22400 = 8.616 C

Now, Q = I × t

I = Q/t 

I = 8.616/1

I = 8.616 ampere

Example 3: How many moles of mercury will be produced by galvanic isolation 1.0 M Hg(NO3)2 solutions with a current of 2.00 A for 3 hours?

Solution:

Hg2+ + 2e– ⇒ Hg

Charged Passed(Q) = I × t(sec)

Q  = 2.0 A × 3.0 × 60 × 60 = 21600 C

2 × 96500 C of Charge produce = 1 mol of Mercury

21600 C of Charge will produce = 1× 21600 / (2×96500) = 0.112mol of Mercury

Example 4: A solution of CuSO4 is electrolyzed for 10 minutes with a current of 1.5 amperes. What is the mass of copper deposited at the cathode?

Solution:

Current (I) = 1.5 A

Time (t) = 10 min = 10 × 60 = 600s 

amount of electricity passed = I × t = (1.5 A) × (600 s) = 900 C (A s = C)

Copper is Deposited: Cu+2 + 2e– → Cu(s)

2 mol of Electrons or 2 × 96500 C of current deposit copper = 63.56 g

900 C of current will deposit copper = 63.56/(2 ×96500) = 0.296 g

Example 5: Calculate how long it will take to deposit 1.0 g of chromium when a current of 1.25 A flows through a solution of chromium (III) sulphate. (Molar mass of Cr=62).

Solution:

Cr3++3e– → Cr(s) 

3 mol of electricity are needed to deposit 1 mol of Cr,

52 g of Cr require current = 3 × 96500 C

1g of Cr will require current = (3 × 96500)/52  = 5567.3 C

Number of Coulombs = Current × t

Time (s) required = No. of Coulombs / Current 

Time (s) required = 5567.3 C / 1.25 (Ampere)

                            = 4453.8s or 1.24hr

Example 6: How many hours does it take to reduce 3 mol of Fe with 2.0 A current? (F= 96500 C) Solution: Reduction of Fe³+ to Fe²+.

Solution:

Fe3+ + e– → Fe²+ 

Reduction of 1 mol of Fe3+ requires = 96500 C

Reduction of 3 mol of Fe³+ require = 3 × 96500 C = 2.895 × 10°C

Quantity of electricity = Current × Time

2.895 × 10 = 2 × Time

Time = 2.895 × 105 

            = 14475 × 105 s

Time = (14475 × 105 )/(60 × 60) 

        = 40.21 hours

FAQs on Electrolysis

Q1: What is Electrolysis?

Answer:

The chemical process that happens when the electric current is passed through an ionic solution is called the electrolysis. In this process, ions move in their aqueous solution forming the compounds on the cathode and anode.

Q2: What is an Electrolytic Cell?

Answer:

A chemical cell with two electrodes a positive electrode called an anode and the negative called the cathode is dipped in the aqueous solution of an ion and electricity is passed through it, which is called an electrolytic cell.

Q3: Where is the Process of Electrolysis used?

Answer:

Electrolysis process is a highly used process in electrochemistry. It is used in the purification of molten metals, breaking an ionic compound into its constituent elements,  electrolysis of water, etc.

Q4: What is Cathode?

Answer:

The electrode of the electrolytic cell which is negatively charged is called the cathode. It attracts the positive cation and the reduction takes place at the cathode.

Q5: What is Anode?

Answer:

The electrode of the electrolytic cell which is positively charged is called the anode. It attracts the positive anion and the oxidation takes place at the anode.



 

Electrochemistry – Cells and Batteries

Last Updated : 14 Feb, 2023

A collection of electrochemical cells used as a power source is referred to as a battery. An oxidation-reduction reaction forms the basis of an electrochemical cell. In general, every battery is a galvanic cell that generates chemical energy through redox reactions between two electrodes. Batteries are globally used in several electronic devices as a source of power.

Working of a Battery 

The battery is an essential component that ensures the smooth operation of many electrical devices. It holds chemical energy and gives various devices electrical energy. The image given below shows us what a conventional cell(battery) looks like.

Cell or Battery

 

The battery’s capacity to work is supported by an electrochemical cell. Electrochemical cells can range in number from one to many in a battery. Two electrodes are present in every electrochemical cell, and an electrolyte separates them. One electrode produces electrons as a result of the chemical process occurring inside the cell. When the electrons start travelling, electricity is created. A chemical process takes place inside a battery, and the electrons move from one electrode to the next to create an electric circuit.

Let’s study battery features and types in the article.

Features of a Battery

A battery should have various features for it to be widely useable. Some of the most important features that a battery should have are,

  • For ease of transportation, it should be small and light.
  • It should last for a respectably longer time both when it is in use and when it is not.
  • A battery or cell must be able to supply a steady voltage. Additionally, the battery or cell’s voltage must not change while being used.

Different Types of Battery

There are primarily two types of batteries or functional cells used commercially.

  • Primary Batteries or Cells
  • Secondary Batteries or Cells

Primary Batteries or Cells

They are known by the name of non-rechargeable batteries. These are the batteries that are only useful when used once. These batteries are not rechargeable or reusable. Alkaline batteries and coin cell batteries are typical examples of primary batteries. Typically, watches, clocks, torches, and other inexpensive electronic gadgets use these types of batteries. These batteries only allow one direction for redox reactions.

Dry Cell

The dry cell, a type of household battery commonly used to power clocks, TV remotes, and other gadgets, is an example of a primary battery. In these cells, a carbon rod serves as the cathode and a zinc container serves as the anode. The cathode is surrounded by a powdered manganese dioxide and carbon combination. A moist paste made of ammonium chloride and zinc chloride is used to fill the area between the container and the rod.

These cells undergo the following redox reaction:

At anode:

Zn(s) → Zn2+ (aq) + 2e–

At cathode:

2e– + 2 NH4+ (aq) → 2 NH3 (g) + H2 (g)

2 NH3 (g) +Zn2+ (aq) → [Zn (NH3)2]2+ (aq)

H2 (g) + 2 MnO2 (S) → Mn2O3 (S) + H2O (l)

Overall cell equation is as follows:

Zn(s) + 2 NH4+ (aq) + 2 MnO2 (S) → [Zn(NH3)2]2+ (aq) + Mn2O3 (S) + H2O (l)

In the area between the cathode and the anode, there is a mixture of MnO2 and a viscous paste of charcoal, zinc chloride, and ammonium chloride (NH4Cl). The porous paper’s lining keeps the paste and zinc container from contacting each other directly. It serves as a bridge for salt. Pitch or wax is used to seal the cell from the top.

Mercury Cell

The mercury cell is a new type of cell that is used in small electrical circuits such as those hearing aids, watches, and cameras. A zinc anode and a mercury (II) oxide cathode make up this component. The electrolyte is a KOH and ZnO paste.

The cell undergoes the following reaction: 

At anode:

Zn(Hg) + 2OH¯→ ZnO(s) + H2O+ 2e¯

At cathode:

HgO(s) + H2O+ 2e¯ → Hg(l) + 2OH¯
 

Overall cell equation is as follows:

Zn + HgO(s) → ZnO(s) + Hg(l)

It has the benefit that its potential stays basically constant during the course of its existence. The mercury cell has a voltage of about 1.35 V.

Secondary Batteries or Cells

These batteries are also called Rechargeable batteries. These batteries are long-lasting, reusable, and excellent for a variety of uses. They are a little more expensive than primary batteries, but when used carefully, safely, and with caution, they last the users longer. Lead-acid batteries and lithium-ion batteries are a few common examples of secondary batteries. The primary applications for these batteries are robots, solar lighting, luxury toys, etc.

Lead Storage Battery

A lead storage battery used in cars and inverters can only be recharged a select number of times. A lead anode and a lead grid filled with lead dioxide make up the cathode of a lead storage battery. As an electrolyte, a 38% concentration of sulfuric acid is utilized.

At anode:

Pb → Pb2++ 2e–

Pb+ SO42– → PbSO4 + 2e–

At cathode:

2e–+ PbO2 + 4H+ → Pb2++ 2H2O

2 e–+ PbO2 + 4H++ SO42- → PbSO4 + 2H2O

These batteries can be recharged by transferring the charge in the other direction and reversing the process, which turns PbSO4 back into Pb and PbO2.

Overall reaction can be written as: 

2PbSO4 (s) + 2H2O → Pb(s) + PbO2 (s) + 2H2SO4

It functions as a voltaic cell and generates electricity when used to start the car’s engine. It functions as an electrolytic cell while being recharged.

Nickel Cadmium Storage Cell

Another rechargeable cell is the nickel-cadmium storage cell. Although it costs more than lead storage batteries, it lasts longer than lead storage cells. However, because it is lighter and smaller, there are certain benefits. Appliances that are portable and cordless can use it.

  • It has a cadmium anode and a metal grid acting as a cathode that contains NiO2.
  • The electrolyte used in this cell is KOH.

The reaction can be written as: 

At anode:

Cd(s) + 2OH¯ → CdO(s) +H₂O(l) + 2e¯

At cathode:

2Ni(OH3)(s) + 2e¯ → 2Ni(OH)2(s) +2OH¯(aq)

Overall cell equation is as follows:

Cd(s) + 2Ni(OH3)(s) → CdO(s) + 2Ni(OH)2(s) + H2O(l)

The reaction byproducts typically attach to the electrodes and can be changed back into something else by charging the cell. Similar to how a lead storage battery is charged.

Lithium Ion Battery

A lithium-ion battery is a specific kind of rechargeable battery that stores energy through the reversible reduction of lithium ions. It is the most common type of battery used in electric vehicles and portable consumer gadgets. Li-ion batteries don’t suffer from the memory effect, have low self-discharge, and have high energy densities.

During a discharge cycle, lithium atoms in the anode are ionized and separated from their electrons. The lithium ions move from the anode through the electrolyte to the cathode, where they combine with their electrons and turn into electrically neutral molecules. Because of their small size, the lithium ions can pass through a micro-permeable separator that separates the anode from the cathode.

Different materials can be used as electrodes in Li-ion batteries. The most common cathode and anode materials are lithium cobalt oxide (cathode) and graphite (anode), and these materials are most commonly found in portable electronic devices like laptops and cell phones. Lithium iron phosphate and lithium manganese oxide, which are used in hybrid and electric vehicles, respectively, are additional cathode components. Ether, a group of organic compounds, is frequently utilized in Li-ion batteries as an electrolyte.

Oxidation-reduction (Redox) reactions take place inside a lithium-ion battery.

Cathode is where reduction takes place. Lithium-cobalt oxide is produced, when cobalt oxide and lithium ions react (LiCoO2). The partial reaction is:

CoO2 + Li+ + e– → LiCoO2

Anode is where oxidation takes place. There, lithium ions and graphite (C6) are formed by the graphite intercalation complex (LiC6). The partial response is:

LiC6 → C6 + Li+ + e–

Complete reaction will be, (right to left = charging; left to right = discharging).

LiC6 + CoO2 ⇄ C6 + LiCoO2

Uses of Battery

Batteries are used for a variety of purposes. The most common uses of batteries are:

  • Batteries are used in Medical Equipment and Home Appliances.
  • Implantable medical devices like pacemakers and insulin pumps utilize bio-batteries.
  • Batteries are used in Construction.
  • Batteries can be used in toys as well as in different gifting products. 
  • Batteries are used in Emergency Response and Firefighting.
  • Batteries are used in Military Operations, surveillance, and spying devices.
  • Batteries are used in Electric and normal automobile batteries. 

Read, More

FAQs on Cells and Batteries

Question 1: What is a Battery?

Answer:

 A device that transforms chemical energy into electrical energy, one or more electrical cells make up a battery.

Question 2: What is the functioning of the dry cell?

Answer:

A wet paste of NH4Cl and ZnCl2 is present in the cell. A dry cell can only function if the paste inside it is moist. Even when the zinc container is not in use, the acidic NH4Cl paste continues to corrode it.

Question 3: What is Electrochemistry?

Answer: 

Electrochemistry is the study of how energy is stored in batteries and changed from one form to another. It is necessary for functioning of the batteries. 

Question 4: What is the life cycle of a battery? 

Answer: 

A battery is considered to have completed a cycle when it is fully charged and then drains to 80% of its initial capacity. The cycle life of a battery is determined by how many cycles it can go through.

Question 5: What are charging current and charging voltage?

Answer: 

Charging Current: It is the highest possible current that can be used to charge the battery.

Charging Voltage: It is the highest possible voltage that should be used to effectively charge the battery.



 

Galvanic Cell

Last Updated : 17 Apr, 2024

Galvanic Cell also called Voltaic Cell is an electrochemical device that converts spontaneous chemical energy generated in a redox reaction into electrical energy.

Table of Content

What is Galvanic Cell?

We define a Galvanic cell as a device that converts the chemical energy of the redox reaction to electrical energy, this is a type of electrochemical cell that uses electrolytes to produce the electrical energy.

To understand the concept of a Galvanic Cell, let’s first understand what is cell and what are its types.

Cell Definition

A Cell is an electrical device that when connected to a circuit generates a potential difference which results in the flow of charge or ions from higher potential to lower potential. A Cell is a unit source of power. When cells are combined together to create potential differences then it is called Battery.

Depending upon the conversion of energy from chemical to electrical or electrical to chemical there are two types of cells:

  • Electrolytic Cell
  • Electrochemical Cell

Electrolytic Cell Definition

An Electrolytic Cell is a device that converts electrical energy into chemical energy. It means it already has the power supply which is used in ‘lysis’ means the breaking of electrolytes into ions which then moves towards electrodes to constitute current and produce electrical energy. In an electrolytic cell, Anode is +ve while Cathode is -ve. In this type of cell, the flow of electrons is from Anode to Cathode.

Learn more about, Electrolytic Cells and Electrolysis

Electrochemical Cell Definition

An Electrochemical Cell is a device that converts chemical energy into electrical energy. It means the chemical energy stored in the cell undergoes a reaction to produce electrical energy. In this type of cell, Anode is -ve while Cathode is +ve. The flow of electrons is from Cathode to Anode. This cell is the reverse of an electrochemical cell.

Let’s learn about Galvanic Cells in detail in this article.

Primary Cell & Secondary Cell

Primary Cells are non-rechargeable cells. Once, the chemical energy stored in the cell consumes the cell becomes useless. It is disposable in nature.

Secondary Cells are rechargeable cells. It is based on the principle of reverse chemical reactions i.e. electrical energy is used to produce electrons inside the cell which is used to run the device. It is economically and environmentally advantageous than Primary Cells.

What is a Galvanic Cell?

The devices in which chemical reaction is used to produce electrical energy are called Galvanic Cells or Voltaic Cells. In these devices, the Gibbs Energy of the spontaneous Redox Reaction is converted into electrical work that can be used to drive a motor or to power electrical equipment such as heaters, fans, geysers, etc.

These cells are greatly important because of their many practical applications. An early example of the Galvanic Cell is Daniel’s Cell invented by the British chemist John Daniels in 1836.

Daniel’s cell was constructed based on the redox reaction:

Zn(s) + Cu²+(aq)+ → Zn²+(aq)+ Cu(s)

Half-Cell Reactions

In these cells, Oxidation and Reduction reactions occur in autonomous containers called half cells or Redox Couples. The half cell of the reaction is represented as follows

Anode(Oxidation): Zn → Zn2+ + 2e–

Cathode(Reduction): Cu2+ + 2e–→Cu

Oxidation occurring at Anode is referred to as Oxidation Half Cell while the Reduction occurring at Cathode is called Reduction Half Cell. Although these half-cell reactions are occurring in separate containers they are connected with each other internally and externally. Internally they are connected via Salt Bridge while externally they are connected via wire, voltmeter, and switch.

The above-shown redox reaction is spontaneous. General Representation of a Galvanic Cell:

M1(s)/M1n+ (aq) || M₂n+(aq)/M₂(s)

Parts of Galvanic Cell

Various parts of the Galvanic Cell include,

  • Anode: The electrode at which oxidation occurs is the anode.
  • Cathode: The electrode at which reduction occurs is the cathode.
  • Salt bridge: It is a tube filled with electrolytes that maintain the neutrality of the Galvanic Cell.
  • Half-cells: Two separate beakers where oxidation and reduction occur are called Half-Cell.
  • External circuit: It helps to conduct the flow of electrons between electrodes.

Constructions of Galvanic Cell

  • A Galvanic Cell is made by combining two electrodes an oxidation electrode and a reduction electrode. Both electrodes individually are called the half cell. 
  • The two half-cells are individually filled with different electrolytic solutions which helps in their particular reaction. Both the half cell are connected to each other using a Salt Bridge internally and externally via wire, switch, and voltmeter.
  • Oxidation occurs at the oxidation electrode which releases the free electrons that accumulate on the electrode and provide a negative potential. The electrode at which oxidation occurs is called the Anode.
  • Reduction occurs at the reduction electrode that generates the positive charge and provides the positive potential. The electrode at which reduction occurs is called the Cathode.
  • Connecting these electrodes via wire, switch, and voltmeter initiates the flow of electrons from one electrode to another resulting in a flow of electric current. For a galvanic cell, the anode is negatively charged and the cathode is positively charged.

A Galvanic cell image is added below in the article.

Principle and Working of Galvanic Cell

The working of Galvanic Cell is discussed below:

  • Electrodes are exposed to the electrolyte at the electrode-electrolyte interface, which generates ions in the electrolyte solution making one metal electrode negatively charged. 
  • For the other electrode, the metal ions in the electrolyte solution deposit on the other metal electrode making the electrode positively charged.
  • Due to this charge separation, a potential difference is developed between electrodes and electrolytes. This potential difference is called electrode potential.
  • The electrode, where oxidation occurs, is the anode while the electrode where reduction occurs is the cathode.
  • Anode is at a negative potential with respect to the solution while the cathode is at a positive potential with respect to the solution.
  • The potential difference between these two electrodes is called the potential of a Galvanic cell and is responsible for the flow of electrons in the circuit.

Electrode Potential

In each half-cell, there is a movement of electrons between the electrodes and the electrolyte. Since there is a flow of charge between the electrode and electrolyte there develops a potential called Electrode Potential. There are two types of Electrode Potential, Oxidation Potential, and Reduction Potential. Their representation is given below:

Oxidation Potential: M → Mn+ + ne–

Reduction Potential: Mn+ + ne– → M

The Electrode Potential is affected by the nature of metal and ion, its concentration, and temperature. The Electrode Potential of a half-Cell is written in terms of its Reduction Potential. Hence, the Electrode Potential of the Oxidation half-cell is -ve and that of the Reduction half-cell is +ve.

Standard Electrode Potential

The Electrode Potential calculated above is relative in nature. In order to find the individual potential of an electrode we use a Standard Hydrogen Electrode whose potential is zero to calculate Standard Electrode Potential.

A Standard Hydrogen Electrode consists of a Platinum Wire covered with Platinum foil in a test tube which is immersed in a 1M concentration of HCl which liberated H+ ion and hydrogen gas is bubbled in it at 1atm at 298K of temperature. 

Standard Electrode Potential is the potential of an Electrode dipped in 1M concentration of its salt in a half cell and this half cell is connected to a Standard Hydrogen Electrode via a salt bridge.

It is represented as

M | Mn+(1M) || H+(1M) | H2 (1 atm), Pt

Standard Electrode Potential

Cell Potential

Cell Potential refers to the potential difference between the cathode and anode of the Galvanic Cell. When no current is drawn from it i.e. the two electrodes are not connected with each other then it is called Cell Electromotive Force or EMF of Galvanic Cell.

The convention to represent cell potential follows that the anode potential is written on the left side and the cathode potential is written on the right side and both are separated by two vertical lines (||). For example, M1(s)/M1n+ (aq) || M₂n+(aq)/M₂(s).

Hence, the left potential is Anode Potential while the right potential is Cathode Potential. Hence, Cell Potential, Ecell is given as

Ecell = Eright – Eleft

Example of Galvanic Cell

Daniel’s cell is the most common example of a galvanic cell. The galvanic cell converts chemical energy into electrical energy. For a Galvanic Cell Copper Ions are reduced at the cathode and Zinc Ions are oxidized at the anode.

Galvanic Cell

Reactions taking place at the cathode and anode of a Galvanic cell are:

At Anode: Zn → Zn2+ + 2e–

At Cathode: Cu2+ + 2e– → Cu

What is Salt Bridge?

Salt Bridge is a U- shaped tube that contains a concentrated solution of inert electrolytes. Some examples of electrolytes used in the salt bridge are KCl, KNO3, K2SO4, etc. These inert electrolytes do not participate in the cell reaction.

Salt Bridge allows the movement of ions from one solution to the other without mixing two solutions. The salt bridge also helps to maintain the electrical neutrality of the solution in the two half-cells.

Difference between Galvanic Cell and Electrolytic Cell

Galvanic Cells and Electrolytic Cells are both electrochemical cells and the major difference between them is as follows:

Galvanic Cell

Electrolytic Cell

It converts chemical energy into electrical energy.It converts electrical energy into chemical energy.
The reactions are spontaneous in Galvanic CellThe reactions are non-spontaneous in Electrolytic Cell.
Both electrodes, cathodes, and anodes are placed in separate beakersBoth electrodes, cathodes, and anodes are placed in the same beaker.
The electrolytes taken in both beakers are different.Only one electrolyte is taken.
Oxidation takes place at the anode (negative end), and reduction takes place at the cathode (positive end).Oxidation takes place at the cathode (positive end), and reduction takes place at the anode (negative end).
A salt bridge is used.No salt bridge is used.
Gibb’s free energy change during the reaction is negative.Gibb’s free energy change during the reaction is positive.

Read More:

Solved Examples on Galvanic Cells

Example 1: Calculate ΔrGφ for the reaction: 

Mg(s)+Cu2+ (aq) →  Mg2 (aq)+Cu(s)

Given E0cell =2.71 V, 1F = 96500 C mol-1

Solution:

ΔrGφ = -nF 
Eocell = 2.71 V,
1 F = 96500 C mol-1, 
n = 2 

ΔrGφ = -2×96500 C mol-1 ×2.71 V

         = -523030 J mol-1                     (1CV = 1J)

         = -523.080 kJ mol-1

Example 2: The  ΔGφ for the Daniell cell has been found to be -212.3 kJ at 25°C. Calculate the equilibrium constant for the cell reaction.

Solution:

ΔGφ =-RT ln Kc

ΔGφ = -212.3 kJ = -212300 J, 
T = 298 K

Here 
R=8.314.K-1 mol-1

ln(Kc) = 212300 / (8.314 × 298) 
          = 85.69

Kc = 1.64 × 1037

Example 3: What does the negative sign in the expression EoZn2+/Zn =-0.76 V mean?

Solution:

It means that zinc is more reactive than hydrogen. When zinc electrode is connected to SHE, zinc will get oxidized and H+ will get reduced.

Example 4: A galvanic cell has an electrical potential of 1.1 V. If an opposing potential of 1.1 V is applied to this cell. What will happen to the reactance of the cell and the current flowing in the cell?

Solution:

When the opposing potential becomes equal to the electric potential, the reaction of the cell stops and no current flows through the cell. Thus, no chemical reaction takes place.

Galvanic Cells – FAQs

What is a Galvanic cell?

An electrochemical cell that converts the chemical energy of redox reactions into electrical energy is called Galvanic Cell or Voltaic Cell.

What is the function of a Galvanic cell?

Galvanic cell is a device which provides Electric energy using Chemical energy. It uses the spontaneous energy of the redox reaction for providing electric energy.

Is Daniel’s Cell a Galvanic cell?

Yes, Daniel’s is a galvanic cell. It is the most common example of a galvanic cell.

How do you make a Galvanic cell?

A galvanic cell is made by dipping two electrodes in a glass vessel solution of dilute sulfuric acid. The two electrodes are made of copper and zinc. The cathode is made of Copper and the anode is made of Zinc.

What is Need for the Salt Bridge in a Galvanic cell?

Salt bridge helps to maintain the neutrality of the solution and allows the free flow of ions from one-half cell to another half cell.

Where does Oxidation Occur in a Galvanic cell?

In a galvanic cell, oxidation occurs at the Anode.

Where does Reduction Occur in a Galvanic cell?

In a galvanic cell, reduction occurs at the Cathode.

What is the Effect of Temperature on the Galvanic Cell?

According to the Nernest Equation, the voltage of the galvanic cell decreases with increasing temperature.

How does a salt bridge function in a Galvanic Cell?

The salt bridge is essential for maintaining electrical neutrality and continuous electron flow in a galvanic cell. It prevents the accumulation of positive and negative ions in the half-cells by allowing ions to migrate, thus completing the electrical circuit and enabling the cell to function effectively​.

What are the key components of a Galvanic Cell?

Key components of a galvanic cell include the anode (where oxidation occurs), the cathode (where reduction occurs), a salt bridge (to maintain ion balance and complete the circuit), electrolytes in each half-cell, and an external wire connecting the electrodes to allow electron flow


 

Fuel Cells – Definition, Types, Advantages, Limitations

Last Updated : 24 Feb, 2022

The study of the link between electrical energy and chemical changes is the subject of electrochemistry, a chemistry subdiscipline. Electrochemical reactions are chemical processes that include the input or creation of electric currents. A fuel cell is an electrochemical cell that uses an electrochemical process to create electrical energy from fuel. To keep the processes that generate electricity going, these cells need a constant supply of fuel and an oxidising agent (usually oxygen). As a result, until the supply of fuel and oxygen is shut off, these cells can continue to generate power.

Fuel cells, while being conceived in 1838, did not enter commercial usage until a century later, when NASA utilised them to power space capsules and satellites. Many establishments, including businesses, commercial buildings, and residential structures, now employ these devices as a major or secondary source of electricity.

Fuel Cell

Fuel cells are cells that directly transform the chemical energy of a fuel cell into electrical energy. Fuels such as hydrogen (H2), carbon dioxide (CO2), methane (CH4), propane (C3H8), methanol (CH3OH), and others are used to create electrical energy in the cells shown below. The fuel cell is constantly supplied with fuel, while the products are continuously removed. There are a great number of fuel cells on the market. The most popular type is a hydrogen-oxygen fuel cell.

Fuel Cell used in the Apollo Space Program

Bacon invented the hydrogen-oxygen fuel cell in 1959. As a result, it’s also called the Bacon cell. It is a possible source of electrical energy that was employed as the major source of electrical energy during the United States’ Apollo space program.

Two porous carbon electrodes are impregnated with a suitable catalyst such as platinum (Pt), silver (Ag), cobalt oxide (CoO), and so on in a basic H2–O2 fuel cell. The electrolyte is a concentrated solution of potassium hydroxide (KOH) or sodium hydroxide (NaOH) that fills the gap between two electrodes. A porous carbon electrode bubbles hydrogen gas (H2) and oxygen gas (O2) into the electrolyte.

  • At the anode, 2H2 (g)+4OH– (aq)  →  4H2O (l)+4e–
  • At the cathode, O2 (g)+2H2O (l)+4e–  →  4OH– (aq)
  • Overall reaction, 2H2 (g) + O2 (g)  →  2H2O (l)

Microbial fuel cells (MFCs) are cells that use microbes to catalyse biological processes to generate power from organic or inorganic chemicals. It would be capable of achieving a 50% efficiency rate. Proteobacteria, Desulfuromonas, Alcaligenes faecalis, Pseudomonas aeruginosa, and other microorganisms have been employed in MFCs.

An anode and cathode compartment are separated by a cation-specific membrane-like potassium nitrate base membrane in a microbial fuel cell. The microorganism oxidises the fuel at the anode, creating carbon dioxide, electrons, and protons. Protons are delivered to the cathode compartment through the membrane and electrons are transferred to the cathode compartment through an external electric circuit. The cathode compartment produces water by mixing oxygen with electrons and protons.

Anode reaction: C12H22O11+13H2O→12CO2+48H++48e–

Cathode reaction: 4H++O2+4e–→2H2O

Types of Fuel Cells

Fuel cells come in a variety of forms.

  1. Polymer Electrolyte Membrane (PEM) Fuel Cell
  2. Phosphoric Acid Fuel Cell
  3. Solid Acid Fuel Cell
  4. Alkaline Fuel Cell
  5. Molten Carbonate Fuel Cell
  6. Hydrogen-Oxygen fuel cell
  7. Microbial fuel cells (MFCs)
  8. Solid Oxide Fuel Cells (SOFCs)
  9. Zinc-Air Fuel Cell (ZAFC)
  10. Direct Methanol Fuel Cell (DMFC)

The Polymer Electrolyte Membrane (PEM) Fuel Cell

  1. Proton exchange membrane fuel cells are another name for these cells (or PEMFCs).
  2. These cells function at temperatures ranging from 50 degrees Celsius to 100 degree Celsius.
  3. The electrolyte used in PEMFCs is a polymer that can conduct protons.
  4. A PEM fuel cell is made up of bipolar plates, a catalyst, electrodes, and a polymer membrane.
  5. Despite its environmentally benign applications in transportation, PEMFCs can also be utilised for fixed and portable power generation.

Phosphoric Acid Fuel Cell

  1. Phosphoric acid is used as an electrolyte in these fuel cells to channel the H+.
  2. These cells operate at temperatures ranging from 150-200 Celsius.
  3. Because phosphoric acid is non-conductive, electrons must go to the cathode via an external connection.
  4. Because the electrolyte is acidic, the components of these cells corrode or oxidise with time.

Solid Acid Fuel Cell

  1. The electrolyte in these fuel cells is a solid acid substance.
  2. At low temperatures, the molecular structures of these solid acids are organised.
  3. At higher temperatures, a phase shift can occur, resulting in a significant increase in conductivity.
  4. CsHSO4 and CsH2PO4 are two examples of solid acids (cesium hydrogen sulphate and cesium dihydrogen phosphate respectively).

Alkaline Fuel Cell

  1. This was the fuel cell that served as the major source of power for the Apollo space programme.
  2. An aqueous alkaline solution is employed in these cells to saturate a porous matrix, which is then used to separate the electrodes.
  3. These cells’ operating temperatures are relatively low.
  4. These cells are quite effective. Along with power, they generate heat and water.

Molten Carbonate Fuel Cell

  1. Lithium potassium carbonate salt is employed as the electrolyte in these cells. At high temperatures, this salt becomes liquid, allowing carbonate ions to migrate.
  2. These fuel cells, like SOFCs, have a relatively high working temperature of 650 Celsius.
  3. Because of the high working temperature and the presence of the carbonate electrolyte, the anode and cathode of this cell are prone to corrosion.
  4. These cells can run on carbon-based fuels like natural gas and biogas.

Solid Oxide Fuel Cells (SOFCs)

Solid oxide fuel cells employ a hard, non-porous ceramic substance as the electrolyte and operate at temperatures between 500 and 1000 degrees Celsius. A solid oxide electrolyte is used in SOFCs to transport negative oxygen ions from the cathode to the anode. SOFCs have an efficiency of 50–60 percent.

  • At the anode: 1/2O2+2e–→O
  • At the cathode: H2+1/2O→H2O+2e–
  • The overall cell reaction: H2+12O2→H2O

Satellites and space capsules employ SOFCs to generate electricity. It is mostly employed in big, high-power applications such as industrial generating plants.

Zinc-Air Fuel Cell (ZAFC)

The Zinc-Air Fuel Cell (ZAFC) is a kind of fuel cell that was created in the United States for use in vehicles. The electrolyte is an aqueous alkali solution such as potassium hydroxide, and the electrode reactions are as follows:

  • Anode: Zn+2OH–→Zn(OH)2+2e–
  • Cathode: O2+2H2O+4e–→4OH–
  • Overall Reaction: 2Zn+O2+2H2O→4Zn(OH)2

It is used as an alternative fuel for vehicles.

Direct Methanol Fuel Cell (DMFC)

Methanol is utilised as a fuel in this subclass of proton-exchange fuel cells. The key benefit of this fuel cell is the ease with which stable liquid fuel methanol may be transported. Polymer membrane serves as the electrolyte, and the electrode reactions are as follows:

  • Anode: CH3OH+H2O→6H++CO2+6e–
  • Cathode: 3/2O2+6H++6e–→3H2O
  • Net reaction: CH3OH+3/2O2→CO2+2H2O

Working of Fuel Cell

A fuel cell may use the chemistry between hydrogen and oxygen to create power. This type of cell was utilised in the Apollo space program and had two purposes: as a source of fuel and as a supply of drinking water (the water vapour produced from the cell, when condensed, was fit for human consumption).

This fuel cell worked by transferring hydrogen and oxygen through carbon electrodes into a concentrated sodium hydroxide solution.

  • Cathode Reaction: O2 + 2H2O + 4e– → 4OH–
  • Anode Reaction: 2H2 + 4OH– → 4H2O + 4e–
  • Net Cell Reaction: 2H2 + O2 → 2H2O

This electrochemical process, however, has a slow response rate. A catalyst, such as platinum or palladium, is used to solve this problem. Before being inserted into the electrodes, the catalyst is finely split to maximise the effective surface area.

Fuel cells have a 70% efficiency in the generation of energy, whereas thermal power plants have a 40% efficiency. Because the creation of electric current in a thermal power plant requires the conversion of water into steam and the use of that steam to move a turbine, there is a significant variation in efficiency. Fuel cells, on the other hand, provide a platform for converting chemical energy into electrical energy directly.

Setup of fuel cells

  • A fuel cell’s primary function is to generate energy, which may be used to power anything from a single light bulb to an entire city. The generation of electricity in a fuel cell is based on a basic chemical reaction that occurs within the cell. The power is then returned to the cell to complete the electric circuit.
  • At the anode, hydrogen atoms are introduced to start the chemical process. At this point, a chemical process removes the electrons from the hydrogen atoms. The hydrogen atoms now have a positive electric charge on them. The cables carry the remaining negatively charged electrons, which create current. At the cathode, oxygen atoms are introduced. They combine with the electrons that the hydrogen atoms have left behind.
  • The oxygen atoms, together with the negatively charged electrons, would either unite with the positively charged hydrogen ions at this point or after passing through the anode, depending on the kind of cell.

Advantages of Fuel Cell

Fuel cells are a possible source of electrical energy, and they offer an advantage over galvanic cells and other traditional techniques of generating electricity by burning fuel. The following are some of the major benefits of fuel cells:

  1. High Efficiency: Fuel cells are theoretically more efficient than traditional techniques for producing electrical energy, such as burning hydrogen, methane, methanol, carbon fuels, or nuclear reactors, since they transfer the energy of a fuel directly into electrical energy. Fuel cells should theoretically be 100% efficient, but only 60–70% efficiency has been achieved thus far. The efficiency of the traditional approach, which involves burning fuel, is only approximately 40%. The thermodynamic efficiency of a fuel cell, n=ΔG/ΔA×100, where, ΔH is the heat of combustion and ΔG is the work done.
  2. Pollution-free Working: The by-products produced by a fuel cell do not pollute the environment. A hydrogen-oxygen fuel cell, for example, generates just water and hence does not contribute to pollution.
  3. Continuous Supply of Energy: As long as fuels are supplied into fuel cells, they can provide energy indefinitely. Unlike traditional cells or batteries, these cells do not experience a decline in voltage or current over time.

Limitations of Fuel Cells

  1. Gaseous fuel is tough to handle. The fuel gas (hydrogen, oxygen, etc.) must be held as a liquid in a specifically built cylinder at a very low temperature and high pressure. This rise is due to the increased cost of the cell, which comes with a number of practical issues.
  2. The catalysts required for electrode reactions, such as platinum (Pt), palladium (Pd), silver (Ag), and others, are highly costly and add to the cell’s cost.
  3. The electrolytes employed in fuel cells are extremely caustic, posing a number of practical issues.

Sample Questions

Question 1: What is a fuel cell?

Answer:

Fuel cells are cells that immediately convert a fuel cell’s chemical energy into electrical energy. In the cells depicted below, hydrogen (H2), carbon dioxide (CO2), methane (CH4), propane (C3H8), methanol (CH3OH), and other fuels are utilised to generate electrical energy. Fuel is continually delivered to the fuel cell, while products are continuously withdrawn. On the market, there are many different types of fuel cells. A hydrogen-oxygen fuel cell is the most common form.

Question 2: What are the types of fuel cells?

Answer:

 Few types of fuel cells are Hydrogen-Oxygen fuel cell, Microbial fuel cells (MFCs), Solid Oxide Fuel Cells (SOFCs), Zinc-Air Fuel Cell (ZAFC), Direct Methanol Fuel Cell (DMFC), etc. 

Question 3: Why do we need fuel cells?

Answer:

Fuel cells are in high demand since they are a cost-effective and ecologically beneficial source of power. Because they may be manufactured individually, they can be utilised for a variety of applications.

Question 4: What are the limitations of fuel cells?

Answer:

  1. Gaseous fuel is difficult to work with. The fuel gas (hydrogen, oxygen, etc.) must be kept as a liquid at a very low temperature and high pressure in a specially designed cylinder. This increase is due to an increase in the cell’s price, which comes with a variety of drawbacks.
  2. The electrode catalysts, such as platinum (Pt), palladium (Pd), silver (Ag), and others, are extremely expensive and add to the cell’s cost.
  3. Fuel cell electrolytes are very caustic, which poses a variety of practical challenges.

Question 5: What are hydrogen fuel cells?

Answer:

The hydrogen fuel cell is a device that directly transforms the chemical energy of hydrogen and oxygen into electricity. Water is created as a by-product of this operation.