Electrochemistry for 12th: Notes for chapter 3 of chemistry 12

Electrochemistry is the most important chapter of physical chemistry. We study about production of electricity from energy released during spontaneous chemical reactions and the use of electrical energy to bring about non-spontaneous chemical transformation. This chapter is important for both purpose as theoretical and practical. As we know that a large number of metals, Sodium hydroxide, chlorine, fluorine and some other chemicals are produced by electrochemical methods. We shall also study in this chapter about various types of fuel cell which convert chemical energy into electrical energy and are used on a large scale in various instruments and devices.

Important contents

• what do you mean by Electrochemistry
• Types of devices used in Electrochemistry
• what is electrolysis and Faraday’s law of electrolysis
• Electrochemical cell and electrode potential
• Determination of electrode potential and Nernst equation.
• Conductance and conductivity and their types
• Kohlrausch law and its application
• Different types of cells and corrosion

What is called Electrochemistry

Eelectrochemistry is the branch of chemistry which deals with conversion of electrical energy into chemical energy or vice-versa with quantitative analysis.

Types of devices used in conversions of energy:

Generally, There are two types of devices are used for conversion of energy from one form to another form.

• Electrolytic cell
• Electrochemical cell

Electrolytic cell : A device that is used to convert electrical energy into chemical energy is called electrolytic cell. Electrochemical cell: A device that is used to convert chemical energy into electrical energy is called electrochemical cell.

Differences between electrochemical and electrolytic cell:

Electrolysis: It is a process in which an electrolyte is taken in molten state or aqueous solution in an electrolytic cell and supplied electricity to decompose into ions and liberated at respective

Faraday’s Law of electrolysis :

Faraday’s 1st law of electrolysis: – According to this law, the liberated mass of any substance in electrolysis is directly proportional to the amount of electricity supplied in electrolytic cell. W ∝ I x t = I x t x Z = Q x Z

• I current in amp,
• t = time in sec,
• Q = quantity of charge (coulomb)
• Z is a constant known as electrochemical equivalent.

When I = 1 amp, t = 1 sec then Q = 1 coulomb, then w = Z. Hence, electrochemical equivalent (Z) is the amount of the substance deposited or liberated by passing of 1A current for 1 sec (i.e.. 1 coulomb, I x t = Q)

Faraday’s 2nd law of electrolysis:

When the same quantity of electricity is passed through different electrolytes. the amounts of the substance deposited or liberated at the electrodes arc directly proportional to their equivalent weights, Hence,

Therefore, electrochemical equivalent ∝ equivalent weight.

Prediction about the Products of Electrolysis:

We can predict about the products of electrolysis in three ways

1. When an electrolyte is taken in molten state and electrodes are of inert metals like Pt, Ni or Pd. The liberated substance corresponds to the ion present in electrolyte. Ex:- when NaCl is allowed to electrolysis, Sodium liberates at cathode and chlorine liberates at anode.
2. When an electrolyte is taken as aqueous solution and electrodes are of inert metal such as Pt or Pd. Liberation of ion depend on discharge potential of ions. Because at the same moment, water also decomposes into H⁺ and OH^– ion.
3. if the cation has higher reduction potential than water (-0.83 V), cation is liberated at the cathode (e.g.. in the electrolysis of copper and silver salts) otherwise H₂ gas is liberated due to reduction of water (e.g., in the electrolysis of K, Na, Ca salts, etc.)
4. Similarly if anion has higher oxidation potential than water (- 1.23 V), anion is liberated (e.g., Br^–), otherwise O₂ gas is liberated due to oxidation of water (e.g., in case of F^–, aqueous solution of Na₂SO₄ as oxidation potential of SO₄^2- is – 0.2 V).
   Discharge potential is defined as the minimum potential that must be applied across the electrodes to bring about the electrolysis and subsequent discharge of the ion on the electrode.
5. When electrodes are common to the metal present in electrolyte, metal reduces at cathode and the ion of the same metal oxidised at anode.

Electrochemicl cell: –

It is a device in which chemical energy is converted into electrical energy spontaneous ly. The set up of this cell can be shown in the following way.

Salt bridge – It is a U – shaped glass tube filled with a paste of KCl, K₂SO₄ or KNO₃ with agar agar. This bridge has following use in electrochemical cell.

1. It completes the circuit and allows the flow of current continuously.
2. It maintains the electrical neutrality on both sides of electrochemical cell. Salt-bridge generally contains paste of strong electrolyte such as KNO₃, KCL etc. Out of them KCI is preferred because the transport numbers of K⁺ and Cl^– are almost same.

Electrode potential- The potential difference between metal and its ion in aqueous solution is called electrode potential. It is denoted by E and its unit is volt(V).

There are two types of electrode potential 1. Oxidation potential (E_ox) in anode cell 2. Reduction potential (E_red) in cathode cell. Or E_ox = – E_red.
Standard electrode potential: If in the half cell, the metal rod is suspended in a solution of one molar concentration, and the temperature is kept at 298K, the electrode potential is called standard electrode potential. it is denoted by E^o . E^o _ox = – E^o_red

Standard hydrogen electrode: It is a reference cell that is used to measure the standard electrode potential of any redox couple of any element. In this electrode, Hydrogen gas at 1 bar pressure is passed into 1M HCl at 298 K in which a foil of platinum coated with platinum black remains immersed. This electrode acts as either as cathode or anode depends on the nature of other

When this electrode acts as cathode, the following reaction takes place

This electrode is usually represented as: Pt, H₂ (g) | H⁺(aq)

Electromotive force or cell potential :

The potential difference between standard reduction potential of two half cell of any electrochemical cell is called cell potential or electromotive force (emf) standard cell potential = E^o cathode – E^oanode Electrochemical series : The arrangement of REDOX couple in increasing order of standard reduction potential is called electrochemical series. In this series, Li has lowest value and F has highest value. Application of electrochemical series:-

• To compares the relative oxidising and reducing powers. Standard reduction potential is directly proportional to oxidising power and inversely proportional to reducing power.
• To compare the relative activities of metals. Standard reduction potential is inversely proportional to the activities of metals.
• To calculate the standard EMF. Where, EMF = E^o cathode – E^o anode
• To predict whether a metal reacts with acid to give hydrogen gas. All the metals having negative standard reduction potential will give hydrogen gas after reaction with protonaed acid.
• To predict the spontaneity of a REDOX reaction. If E_cell value will be +ve, the process is spontaneous.

Nernst eqation :

This equation is applied to measure the electrode potential or cell potential of a redox couple at any concentration or pressure and temperature. This equation can be shown as: For an electrode:

For an electrochemical cell:

♡ – Concentration of pure solid and liquid is taken as unity. Relation between Nernst equation and Equilibrium constant (K)

Concentration cell: The electrochemical cell in which electrodes and electrolytic solutions are same but concentration of solutions are different is called concentration cells.

Molar Conductivity (Λ_m)

The conductivity of all the ions produced when 1 mole of an electrolyte is dissolved in V mL of solution is known as molar conductivity.

It is related to specific conductance as

Λ_m = (k x 1000/M)

where. M = molarity.

It units is Ω^-1 cm² mol^-1 or S cm² mol^-1.

Equivalent conductivity (Λ_m)

The conducting power of all the ions produced when 1 g-equivalent of an electrolyte is dissolved in V mL of solution, is called equivalent conductivity. It is related to specific conductance as

Λ_m = (k x 1000/N)

where. N = normality.

Its units ohm^-1 cm² (equiv^-1) or mho cm² (equiv^-1) or S cm² (g-equiv^-1).

Debye-Huckel Onsagar equation It gives a relation between molar conductivity, Λ_m at a particular concentration and molar conductivity Λ_m at infinite dilution.

Λ_m = Λ⁰_m – √C

where, b is a constant. It depends upon the nature of solvent and temperature.

Factors Affecting Conductivity

(i) Nature of electrolyte The strong electrolytes like KNO₃ KCl. NaOH. etc. are completely ionised in aqueous solution and have high values of conductivity (molar as well as equivalent).

The weak electrolytes are ionised to a lesser extent in aqueous solution and have lower values of conductivity (molar as well as equivalent) .

ii) Concentration of the solution The concentrated solutions of strong electrolytes have SIgnificant interionic attractions. which reduce the speed of ions and lower the value of Λ_m. and Λ_eq.

The dilution decreases such attractions and increase the value of Λ_m and Λ_eq.

The limiting value, Λ⁰_m or Λ^∞_m. (the molar conductivity at zero concentration (or at infinite dilution) can be obtained extrapolating the graph.

In case of weak electrolytes, the degree of ionization increases dilution which increases the value of Λ_m and Λ_eq. The limiting value Λ⁰_m cannot be obtained by extrapolating the graph. ~

limiting value, Λ⁰_m, for weak electrolytes is obtained by Kohlrausch law.

(iii) Temperature The increase of temperature decreases inter-ionic attractions and increases kinetic energy of ions and their speed. Thus, Λ_m and Λ_eq increase with temperature.

Kohlrausch’s Law

At infinite dilution, the molar conductivity of an electrolyte is the sum of the ionic conductivities of the cations and anions, e.g., for A_xB_y.

Applications of Kohlraush Law

(i) Determination of equivalent/molar conductivities of weak electrolytes at infinite dilution, e.g.,

(ii) Determination of degree of dissociation (α) of an electrolyte at a given dilution.

The dissociation constant (K) of the weak electrolyte at concentration C of the solution can be calculated by using the formula

k_c = (Cα²/1 – α)

where, α is the degree of dissociation of the electrolyte.

(iii) Salts like BaSO₄ .., PbSO₄‘ AgCl, AgBr and AgI which do not dissolve to a large extent in water are called sparingly soluble salts. The solubility of a sparingly soluble salt can be calculated as

Batteries

These are source of electrical energy which may have one or more cells connected in series. For a good quality battery it should be reasonably light. compact and its voltage should not vary appreciably during its use.

Primary Batteries

In the primary batteries. the reaction occurs only once and after use over a period of time battery becomes dead and cannot be reused again.

(i) Dry cell or Leclanehe cell

Anode-Zinc container

Cathode-Graphite rod surrounded by MnO₂ powder

Electrolyte-Paste of NH₄Cl + ZnCl₂

Cathode reaction,

2MnO₂(s) + 2 NH⁺₄(aq) + 2e^– → Mn₂O³(s) + 2NH₃(g) + H₂O(l)

Anode reaction,

Zn(s) → Zn²⁺(aq) + 2e^–

Cell potential 1.25 V to 1.5 V

(ii) Mercury cell

Anode-Zn-Hg amalgam

Cathode-Paste of (HgO + C)

Electrolyte-Moist paste of KOH-ZnO

Secondary Batteries

These cells can be recharged and can be used again and again, e.g.,

(i) Lead Storage battery

Anode-Spongy lead

Cathode-Grid of lead packed with PbO₂

Electrolyte-38% H₂SO₄ by mass

When recharged the cell reactions are reversed.

(ii) Nickel-cadmium storage cell

Anode-Cadmium

Cathode-Metal grid containing NiO₂

Electrolyte-KOH solution

Anode reaction,

Cd(s) + 2OH^–(aq) → Cd(OH)₂(s) + 2e^–

Fuel Cells

Galvanic cells which use energy of combustion of fuels like H₂, CH₄, CH₃OH, etc., as the source to produce electrical energy are called fuel cells. The fuel cells are pollution free and have high efficiency.

Hydrogen-Oxygen Fuel Cell

Electrodes-Made of porous graphite impregnated with catalyst (Pt, Ag or a metal oxide).

Electrolyte-Aqueous solution of KOH or NaOH

Oxygen and hydrogen are continuously fed into the cell.

Corrosion

Slow formation of undesirable compounds such as oxides, sulphides or carbonates at the surface of metals by reaction with moisture and other atmospheric gases is known as corrosion.

Factors Affecting Corrosion

1. Reactivity of metals
2. Presence of moisture and atmospheric gases like CO₂, SO₂, etc.
3. Presence of impurities
4. Strains in the metal
5. Presence of electrolyte

Rusting of Iron-Electrochemical Theory

An electrochemical cell, also known as corrosion cell, is developed at the surface of iron.

Anode- Pure iron

Cathode-Impure surface

Rusting of iron can be prevented by the following methods :

1. Barrier protection through coating of paints or electroplating.
2. Through galvanisation or coating of surface with tin metal.
3. By the use of antirust solutions (bis phenol).
4.  By cathodic protection in which a metal is protected from corrosion by connecting it to another metal that is more easily oxidised.

Conclusion in Electrochemistry

Electrochemistry is the branch of chemistry in which we learn about conversion of chemical energy into electrical energy and vice-versa. Electrochemical cell is used to convert chemical energy into electrical energy. Electrolytic cell is used to convert electrical energy into chemical energy for electrolysis. Faraday’s law is very important to calculate the required energy for proper yield. Nernst equation is very useful to calculate the electrode potential and cell potential of redox couple.

Conductivity, molar conductivity, equivalent conductivity are important terms which help us to find conductivity of electrolytes in different concentrations. Limiting molar conductivity of any electrolyte is the sum of the limiting Ionic conductivity of each ions produced by ionisation of 1 mole of an electrolyte.

Kohlrausch law is used to determine the limiting molar conductivity of a weak electrolyte. Dry cells, mercury cells are primary cells. Lead storage battery and nickel – cadmium cells are secondary cells and H₂ – O₂ is fuel cell. Rusting of iron is a good example of corrosion which behaves like electrochemical cell.

FAQS in Electrochemistry