# Electrochemistry Class 12 Notes Chemistry Chapter 3

Electrochemistry, Introduction, Electrolytic Conductance, Kohlrausch’s Law, Nernst Equation, Faraday's Laws of Electrolysis, Electrochemical cell

## Introduction

In this chapter, we will begin with the electrical properties of conductors and their types. Then we will proceed towards electrochemical cell and electrolytic cell in which we will study about their functions, applications and relationship with temperature. Then we will move towards types of electrode potentials, batteries and corrosion. So, let us experience an electrifying tour, named "Electrochemistry".

## Electrochemistry

Electrochemistry is the branch of chemistry which establishes a relationship between chemical energy and electrical energy. It deals with study of chemical changes which occur on passing electric current into certain chemical systems and also with the generation of electricity by carrying chemical reactions. The chemical changes which involve the flow of electric current are called electrochemical changes.

## Electrolytic Conductance

All those substances which may allow the passage of electricity through them are called conductors. They are of two types:

### i). Metallic conductors

Conductance in metals is due to mobility of electrons (electron sea-model). Metal may conduct electricity in solid state, molten state or even in dissolved state. With rise in temperature magnitude of metallic conductance decreases.

### ii). Electrolytic conductors (electrolytes)

Conductance is due to movement of ions towards oppositely charged electrodes. It causes transfer of matter and hence decomposition.

• Electrolytes which are almost completely ionized are strong electrolytes. For example, NaCl, KCl, HCl, NaOH, NH4NO3 etc.

• Electrolytes which do not completely ionize in solution are weak electrolytes. For example, CH3COOH, H2CO3, H3BO3, ZnCl2, HgCl2, HCN, NH4OH etc.

#### a). Ohm’s Law

The current (I) carried by a conductor (or electrolytic solution) is directly proportional to the potential difference (V) between the two ends of conductor.

V ∝ I ⇒ V = RI

#### b). Resistance (R)

The constant of proportionality in Ohm’s law is called resistance offered by conductor. Its SI unit is Ω (ohm) or (volt/ampere).

R=\frac{V}{I}

Resistance of a conductor is directly proportional to its length (l) and inversely proportional to its area of cross-section (A)

R∝\frac{l}{A} ⇒ R=ρ\frac{l}{A}

where, ρ = specific resistance or resistivity.

#### c). Specific resistance (resistivity) (ρ)

Resistance offered by a conductor of unit length and unit cross-sectional area is called specific resistance.

ρ=R\frac{A}{l}=\frac{R\times 1}{1}=R

#### d). Conductance (G)

It is the property of conductor (metallic as well as electrolytic) which facilitates the flow of electricity through it. It is equal to reciprocal of resistance. Its SI unit is Ohm–1–1) or mho.

G=\frac{1}{R}

#### e). Specific Conductance (conductivity) κ

It is reciprocal of specific resistance.

κ=\frac{1}{ρ}=\frac{l}{R.A}=\frac{1}{R}.\frac{l}{A}

#### f). Cell Constant

For any cell, the ratio of distance between electrodes and area of electrode is constant, called cell constant. It is denoted by G*.

G* = (\frac{l}{A})

#### g). Molar Conductance (ΛM)

It is defined as conductance of all the ions produced by ionization of one molar solution of an electrolyte. The SI unit of molar conductance is S m2 mol–1.

Λ_{M}=\frac{1000\times κ}{M}

#### h). Equivalent Conductance (λeq)

It is defined as conductance of all the ions produced by ionization of one normal solution of electrolyte. The SI unit of equivalent conductance is S m2 equiv-1.

Λ_{eq}=\frac{1000\times κ}{N}

Also, Λ_{M} =Λ_{eq} × n-factor

## Kohlrausch’s Law

At infinite dilution, an ionic species (cation or anion) contributes a fixed value at a given temperature, towards total conductance of the electrolyte, irrespective of other ionic species in combination with it.

If, λ_{A}^{o} and λ_{B}^{o} be limiting molar conductivities of cation and anion respectively then conductance of the electrolyte is given by

Λ_{A B}^{o}=xΛ_{+}^{o} + yΛ_{-}^{o}

### (i). First Law

Amount of a substance deposited or liberated at any electrode is directly proportional to the charge passed.

W ∝ Q

where, Q is amount of charge flowing in the circuit.

W = Z ⋅ Q or W = Z × i × t

where, i = current

t = time

Z = electrochemical equivalent

### (ii). Second Law

When the same quantity of electricity is passed through solutions of different electrolytes connected in series, the weight of the substances produced at the electrodes are directly proportional to their equivalent weights.

\frac{m_1}{m_2}=\frac{E_1}{E_2}

\frac{Z_{1}it}{Z_{2}it}=\frac{E_1}{E_2}

\frac{Z_1}{Z_2}=\frac{E_1}{E_2}

E∝Z

E=FZ

where, F = farade constant = 96500 C/mol

## Electrochemical Cell

It is a device which converts chemical energy into electrical energy. Redox reaction is responsible for the production of electricity. Oxidation and reduction takes place in two separate half cells.

This cell contains two half cells, one contains ZnSO4 and the other contains CuSO4 solution. The two half cells are connected by an inverted U-tube which is known as salt bridge. It contains the concentrated solution of inert electrolyte like KCl, KNO3, NH4NO3 etc. For the preparation of salt bridge, agar-agar or gelatin is dissolved in hot concentrated aqueous solution of inert electrolyte.

### Working of Cell

When the zinc and copper electrodes are joined by a wire, flow of electrons take place through an external circuit. The direction of flow of electrons is from zinc to copper electrode. In first half cell, oxidation takes place so that Zn+2 are formed and two electrons enter the external circuit. In second half cell, reduction takes place so that copper is formed by taking two electrons from external circuit. Solutions of both compartments remain electrically neutral.

### i). Cell-Representation

Zn | Zn^{+2}(C_{1})|| Cu^{+2}(C_{2})| Cu

Anode reaction

Zn ⟶ Zn+2 (C1) + 2e (oxidation)

Cathode reaction

Cu+2 (C2) + 2e ⟶ Cu (reduction)

Cell reaction (redox reaction)

Zn + Cu+2 (C2) ⟶ Zn+2 (C1) + Cu

### ii). EMF of a Cell

When an electrode is kept in contact with a solution of its ions, the metal tend to lose electrons and thus passes into the solution in the form of metal ions. Thus, an electrical potential difference is set up between metal and its solution. This is known as half cell electrode potential. The difference between the electrode potential of two electrodes constituting an electrochemical cell is called electromotive force (EMF) of cell.

E_{cell}^{0}=E_{Right}^{0}–E_{Left}^{0}

### iii). Standard Electrode Potential

When a metal is placed in 1 M solution of its ions at 25°C then a potential difference develops between metal and its solution. This potential difference is known as standard electrode potential.

### iv). Standard Hydrogen Electrode

The electrodes potential of an electrode can be determined by connecting this half cell with a standard hydrogen electrode. The electrode potential of the standard hydrogen electrode is taken as zero. The electrode potential of a metal electrode as determined with respect to a standard or normal hydrogen electrode is called standard electrode potential (E°).

## Electrochemical Series

When electrodes in contact with their ions are arranged in an increasing order of their standard reduction potential, the resulting series is known as electrochemical series.

### Application of Electrochemical Series

1. The metal which has low standard reduction potential (S.R.P.) is more electropositive.

2. The metals which have low S.R.P. value are highly reactive.

(i). Alkali metal and alkaline earth metal have low S.R.P. value, so they are highly reactive and evolve H2 from cold water.

(ii). Moderately electropositive metal are less reactive and evolve H2 from steam.

(iii). Weakly electropositive metal are not able to evolve H2 from water.

3. The metal which has low S.R.P. (more negative) value can replace other metal from its salt.

4. Metals which have low S.R.P. (more negative) value are strong reducing agents.

5. Non-metals which have higher S.R.P. (more positive) value and are strong oxidising agents.

6. Metals which have low S.R.P. value, their oxides are thermally stable. Metal which have high S.R.P. value, their oxides are thermally unstable.

## Nernst Equation

Let us consider cell reaction

aA + bB ⟶ cC + dD

E_{cell}=E_{cell}^{°}-\frac{2.303 RT}{nF}log\frac{[C]^{c}[D]^{d}}{[A]^{a}[B]^{b}}

E = emf of cell, n = Number of e exchange, F = Faraday constant

At room temperature,

E_{cell}=E_{cell}^{°}-\frac{2.303\times 8.31\times 298}{n\times 96500}log\frac{[C]^{c}[D]^{d}}{[A]^{a}[B]^{b}}

E_{cell}=E_{cell}^{°}-\frac{0.0591}{n}log\frac{[C]^{c}[D]^{d}}{[A]^{a}[B]^{b}}

This equation is known as Nernst's equation.

## Batteries

Two or more electrochemical cells connected in series constitute a battery. These are the following types of commercial cells.

(i) Primary Cells: Those cells which cannot be recharged are known as primary cells. e.g., Mercury cell, dry cell.

(ii) Secondary Cells: Secondary cells can be recharged by passing a current. e.g., Ni – Cd cell, Lead-storage battery.

Six cells are connected in series, each cell provided 2 V, so total voltage provided by the battery is 12 V. The anode, a series of lead grids packed with spongy lead and cathode, a series of grids packed with lead dioxide 38% by wt. H2SO4 acts as an electrolyte. Cell reaction (when discharging take place

Anode :  Pb(s) + HSO4(aq) ⟶ PbSO4(s) + H+ + 2e,  E° = 0.296 V

Cathode :  PbO2(g) + 3H+(aq) + HSO4(aq) + 2e ⟶ 2PbSO4(s) + 2H2O(l),  E° = 1.628 V

Overall :  Pb(s) + PbO2(s) + 2H+ (aq)+ 2HSO4 (aq) ⟶ 2PbSO4(s) + 2H2O(l),  E° = 1.924 V

## Corrosion

Process of slowly eating away of the metal due to attack of atmospheric gases, on the surface of the metal resulting into the formation of oxides, sulphides, carbonates is called corrosion. Corrosion of iron is called rusting. Rust is hydrated ferric oxide Fe2O3 . xH2O.

### Factors which Promote Corrosion

• Reactivity of metal
• presence of impurity
• presence of air and moisture
• strains in metals
• presence of electrolyte.

### Prevention of Corrosion

1. Barrier protection: By using paints, thin film of oil, grease, Cu, Sn etc. If the coating is broken, iron will get rusted.

2. Sacrificial Protection: Covering the surface of iron with a layer of metal which is more active than iron prevents the iron from losing electrons.

Galvanisation: Covering iron with metal zinc. The layer of Zn on the iron surface when comes in contact with moisture, oxygen and CO2 in air, protective invisible thin layer of basic zinc carbonate ZnCO3. Zn(OH)2 is formed due to which the galvanised iron sheets lose their lusture and tends to protect it from further corrosion.

3. Electrical Protection or Cathodic protection: More electropositive metals like Zn, Mg or Al may be connected with the iron pipes burried in the moist soil, canals, storage tanks, etc. Iron will act as cathode and will not get rusted.

4. Using Anti Rust Solutions: These are alkaline phosphate and alkaline chromate solutions. The alkaline nature of solutions prevent availability of H+ ions. On the surface of iron, a protective, insoluble thin film of iron phosphate is formed. These are used in car radiators to prevent rusting.

## Summary

1. In electrochemical cell, the chemical energy is converted into electrical energy.

2. In electrolytic cell the electrical energy is converted into chemical energy.

3. The reduction of cations is based on the standard reduction potential provided all ions have 1 M concentration.

4. From anions the oxidation is based on standard oxidation potentials provided they are at 1 M concentration.

5. For non redox reaction EMF is not defined.

6. Concentration cells are those whose E0cell is zero. They are categorized as electrode and electrolyte concentration cells.

7. Lead storage battery is reversible cell functioning as electrochemical cell during discharging and as electrolytic cell during charging.