ELECTROCHEMISTRY

PREPARED

BY

GEETHA K

PGT CHEMISTRY, JNV KANNUR , KERALA

Why Study Electrochemistry?

It is important because

• Many of devices that we

use every day that are

battery powered.

• A big problem called corrosion

is an electrochemical phenomenon.

• Many chemicals and elements

such as Cl₂, NaOH, Al etc.

are produced electrolytically.

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DEFINITION

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It is the study of the inter-conversion between chemical energy and electrical energy.

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ELECTROCHEMICAL CELLS

What is an Electrochemical Cell?

An electrochemical cell is a device that can generate electrical energy from the chemical reactions occurring in it, or use the electrical energy supplied to it to facilitate chemical reactions in it. These devices are capable of converting chemical energy into electrical energy, or vice versa.

Such cells capable of generating an electric current from the chemical reactions occurring in them care called Galvanic cells or Voltaic cells. Alternatively, the cells which cause chemical reactions to occur in them when an electric current is passed through them are called electrolytic cells.

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GALVANIC CELL

ELECTRO CHEMICAL CELLS�A COMPARISON

DANIEL CELL

In Daniel’s cell, copper ions are reduced at the cathode while zinc is oxidized at the anode.

Reactions of Daniel cell at cathode and anode are:

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At cathode: Cu ²⁺ + 2e^– → Cu

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At anode: Zn → Zn²⁺ + 2e^–

ELECTRODE POTENTIAL

 Following two changes occur when a metal rod is dipped in its salt solution, 

(a) Oxidation: Metal ions pass from the electrode into solution leaving an excess of electrons and thus a negative charge on the electrode.

The conversion of metal atoms into metal ions by the attractive force of polar water molecules.

M →  Mⁿ + ne^-

(b) Reduction: Metal ions in solution gain electrons from the electrode leaving a positive charge on the electrode. Metal ions start depositing on the metal surface leading to a positive charge on the metal.

Mⁿ⁺ +  ne^- →  M

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In the beginning, both these changes occur with different speeds but soon an equilibrium is established.

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M Mⁿ⁺ + ne^-

In practice, one effect is greater than the other,

If first effect is greater than the second, the metal acquires a negative charge with respect to solution and

If the second is greater than the first, it acquires positive charge with respect to solution, thus in both the cases a potential difference is set up.

REFERENCE ELECTRODE

• A reference electrode is an electrode which has a stable and well-known electrode potential.
• There are many ways reference electrodes are used. The simplest is when the reference electrode is used as a half-cell to build an electrochemical cell. This allows the potential of the other half cell to be determined. An accurate and practical method to measure an electrode's potential in isolation (absolute electrode potential) has yet to be developed.

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STANDARD HYDROGEN ELECTRODE (SHE)

• It is reference electrode consists of a platinum electrode in contact with H₂ gas (1 atm) and aqueous H⁺ ions (1 M).
• It is assigned 0.0 V electrode potential.
• It may behave as anodic or cathodic half cell.
• It is represented as

Pt(s)|H₂(g)(a_H2= 1)|H⁺(aq)(a_H+ = 1).

• When SHE is coupled with an other half cell then cell potential is the value of the electrode potential of half cell.

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Determination of Standard Electrode Potential of Zn/Zn²⁺ Electrode�

• A zinc rod is dipped in 1 M zinc sulphate solution. This half-cell is combined with a standard hydrogen electrode through a salt bridge.
• Both the electrodes are con­nected with a voltmeter.
• The deflection of the voltmeter indicates that current is flowing from hydrogen electrode to metal electrode or the electrons are moving from zinc rod to hydrogen electrode.
• The zinc electrode acts as an anode and the hydrogen electrode as cathode and the cell can be represented as     

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• The EMF of the cell is 0.76 volt

E⁰_{CELL =}E⁰_CATHODE - E⁰_ANODE

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0.76 =  E^o_Anode + 0

E^o_Anode = +0.76 V

NERNST EQUATION

Reduction Potential under Non-standard Conditions is determined using Nernst Equation when Concentrations not-equal to 1M.

where:

E = actual ½ cell reduction potential

E^o = standard ½ cell reduction potential

n = number of electrons in reaction

T = temperature (K)

R = ideal gas constant (8.314J/(K-mol)

F = Faraday’s constant (96500 C/mol)

Nernst Equation for a cell

Reduction Potential under Non-standard Conditions is determined using Nernst Equation when Concentrations not-equal to 1M. Thus For the cell,

Cu(s)I Cu²⁺(aq)IIZn²⁺ (aq)IZn(s)

With cell reaction. Cu²⁺ (aq) + Zn(s) Zn²⁺ (aq) +Cu(s)

NERNST EQUATION......

At equilibrium E_cell =0:

The value of E˚_cell is related to Gibbs free energy, ∆G˚ by:

∆G˚ = –nFE˚_cell

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The value of Equilibrium constant, K is related to ∆G˚ by:

∆G˚ = –2.303RT logK

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EQUILIBRIUM CONSTANT AND GIBB’S FREE ENERGY FROM NERNST EQUATION

At equilibrium Ecell = 0

There fore E⁰cell = 0.059/n log Kc

log Kc= E⁰cellX n/0.059

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Δ⁰rG = – nFE⁰⁽cell)

Where Δ⁰rG is standard Gibb’s free energy of reaction and E⁰⁽cell) is standard emf of the cell

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RESISTENCE & CONDUCTIVITY

• Resistance, R, is proportional to the distance, l, between the electrodes and is inversely proportional to the cross-sectional area of the sample, A (noted S on the Figure above). Writing ρ (rho) for the specific resistance (or resistivity),

R =l/Aρ

• In practice the conductivity cell is calibrated by using solutions of known specific resistance, ρ^*, so the quantities l and A need not be known precisely. If the resistance of the calibration solution is R^*, a cell-constant, C, is derived

R*= C X ρ*

• The specific conductance (conductivity), κ (kappa) is the reciprocal of the specific resistance.

κ = 1/ρ = C/R

• Conductivity is also temperature-dependent. Sometimes the ratio of l and A is called as the cell constant, denoted as G^*, and conductance is denoted as G. Then the specific conductance κ (kappa), can be more conveniently written as

κ = G* X G

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MOLAR CONDUCTIVITY

• The conductivity is defined as the conductivity of an electrolyte solution divided by molar concentration
• Molar conductivity = λ_m = κ/c
• if k is expressed in Sm^–1 and the concentration, c in mol m^–3 then the unit of λ_m will be Sm²mol^–1.
• If we use Scm^–1 as the units for k and mol cm^–3, the units of concentration, then the units for λ_m are Scm²mol^–1. It can be calculated by using the equation:
• λ_m (Scm²mol^–1) = κ(S cm^–1) × 1000 (cm³/L)/molarity (mol/L).
• OR, 1 Sm²mol^–1 = 10⁴ Scm²mol^–1

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CONDUCTIVITY CELL

• While measuring the resistance (DC) changes the composition of the solution. So we use an alternating current (AC) source.

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• A solution cannot be connected to the bridge like a metallic wire, so we use a specially designed vessel called conductivity cell.

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VARIATION OF CONDUCTIVITY & MOLAR CONDUCTIVITY WITH DILUTION

• Conductivity decreases with dilution
• Molar conductivity increases with dilution

For strong electrolytes

Λm=Λ⁰m-AC^1/2

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KOHLRAUSCH’S LAW

• It states that limiting molar conductivity of an electrolyte is the sum of the individual contributions of the anion and cation of the electrolyte.

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• Thus, if λ°_Na+ and λ°_{Cl -} are limiting molar conductivity of the sodium and chloride ions , then the limiting molar conductivity for sodium chloride is given by the equation:

Ë°_m(NaCl) = λ°_Na+ + λ°_{Cl - .}

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APPLICATION OF KOHLRAUSCH’S LAW

• This law may be used to determine the limiting molar conductivity, ‘λ°_m’ degree of dissociation ‘α’ and dissociation constant ‘K_a’ of a weak electrolyte.

ELECTROLYTIC CELL AND ELECTROLYSIS

• Electrolysis: It is the process in which electrical energy is used to drive a non-spontaneous chemical reaction.

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• An electrolytic cell is an apparatus for carrying out electrolysis.

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• Processes in an electrolytic cell are the reverse of those in a galvanic cell.
• Electrolysis process is used in Manufacture of Cl₂ and NaOH, Electro-refining and Electroplating, Electrolysis of water etc.

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FARADAY’S LAW OF ELECTROLYSIS

• (i) First Law: The amount of chemical reaction which occurs at any electrode during electrolysis by a current is proportional to the quantity of electricity passed through the electrolyte (solution or melt).
• (ii)Second Law: The amounts of different substances liberated by the same quantity of electricity passing through the electrolytic solution are proportional to their chemical equivalent weights (Atomic Mass of Metal ÷ Number of electrons required to reduce the cation).

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BATTERIES

• Batteries are the most important practical application of galvanic cells.
• Single-cell batteries consist of one galvanic cell.
• Multi-cell batteries consist of several galvanic cells linked in series to obtain the desired voltage.

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Types of batteries

PRIMARY BATTERY

• In these batteries, the reaction occurs only once and after use over a time period battery becomes dead and cannot be reused again.
• Example: Dry Cell

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SECONDARY BATTERY

• A secondary battery after use can be reused by recharging by passing current through it in the opposite direction
• Example: Lead storage cell

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DRY CELL

A dry cell is a type of electric battery, commonly used for portable electrical devices. It was developed in 1886 by the German scientist Carl Gassner, after development of wet zinc-carbon batteries by Georges Leclanché in 1866. The modern version was developed by Japanese Yai Sakizo in 1887.

A dry cell uses a paste electrolyte, with only enough moisture to allow current to flow

A common dry cell is the zinc-carbon cell, sometimes called the dry Leclanché cell, with a nominal voltage of 1.5 volts, the same as the alkaline cell.

• It is also called Leclanche cell
• Anode: Zinc metal can
• Cathode: MnO₂ and carbon paste
• Electrolyte: NH₄Cl and ZnCl₂ paste.
• Cell Potential: 1.5 V but decreases to 0.8 V with use.

Anode: Zn(s) → Zn²⁺ (aq) + 2e^-

Cathode: 2NH⁺₄ (aq) + 2MnO₂ (s) + 2e^- → Mn₂O₃ (s) + 2NH₃ (aq) + H₂O (l)

Overall: Zn(s) + 2NH⁺₄(aq) + 2MnO₂(s) → Zn²⁺ (aq) + 2NH₃(aq) + H₂O(l) + Mn₂O₃(s)

MERCURY CELL/ BUTTON CELL

A mercury battery (also called mercuric oxide battery, mercury cell, button cell, or Ruben-Mallory) is a non-rechargeable electrochemical battery, a primary cell.

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Mercury batteries use a reaction between mercuric oxide and zinc electrodes in an alkaline electrolyte.

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The voltage during discharge remains practically constant at 1.35 volts, and the capacity is much greater than that of a similarly sized zinc carbon battery. Mercury batteries were used in the shape of button cells for watches, hearing aids, cameras and calculators, and in larger forms for other applications

Anode: Zn(Hg) + 2OH^- (aq) → ZnO(s) + H₂O (l) + 2e^-��Cathode: HgO(s) + H₂O (l) + 2e^- → Hg(l) + 2OH^- (aq)��Overall: Zn(Hg) + HgO(s) → ZnO(s) + Hg (l)

Anode : Pb(s) + SO^2- (aq) PbSO₄(s) + 2e^-

Cathode:

PbO₂(s) + 4H⁺ (aq) + SO₄^2- (aq) + 2e^- PbSO₄(s) + 2H₂O(l)

Overall reaction:

Pb(s) + PbO₂(s) + 4H⁺(aq) + 2SO₄^2-(aq) 2PbSO₄ (s) + 2H₂O (l)

LEAD STORAGE CELL

NICKEL CADMIUM BATTERY

• The nickel–cadmium battery (Ni-Cd battery or NiCad battery) is a type of rechargeable battery using nickel oxide hydroxide and metallic cadmium as electrodes.

A fully charged Ni-Cd cell contains:

• a nickel(III) oxide-hydroxide positive electrode plate
• a cadmium negative electrode plate
• an alkaline electrolyte (potassium hydroxide).
• anode reaction
• (i) Cd_(s) + 2OH^-_(aq) ==> Cd(OH)_2(s) + 2e^-

cathode reaction

• C

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• It is rechargeable
• Cell Potential :1.30 V
• Electrolyte: NiO(OH).

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Anode: Cd(s) + 2OH^–(aq) → Cd(OH)₂(s) + 2e^–

cathode reaction

(ii) NiO_2(s) + 2e^-  + 2H₂O ==> Ni(OH)_2(s) + 2OH^-_(aq)

overall cell reaction

(iii) Cd_(s) + NiO_2(s) + 2H₂O ==> Cd(OH)_2(s) + Ni(OH)_2(s)

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FUEL CELL

• It’s a type of galvanic cell that requires a continuous supply of reactants to keep functioning.
• Fuel cells are not batteries because they are not self-contained.
• It uses externally fed CH₄ , CH₃OH or H₂, which react to form water.
• Electrolyte: Hot aqueous KOH solution.
• Cell Potential: 1.23 V and have about 40% conversion to electricity; the remainder is lost as heat. Excess heat can be used to drive turbine generators.

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Anode : 2H₂ (g) + 4OH^- (aq) 4H₂O (l) + 4e^-

Cathode: O₂ (g) + 2H₂O (l) + 4e^- 4OH^- (aq)

Overall: 2H₂ (g) + O₂ (g) 2H₂O (l)

CORROSION

• Corrosion is the oxidative deterioration of metal.

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• 25% of steel produced goes to replace steel structures and products destroyed by corrosion.

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• Rusting of iron requires the presence of both oxygen and water.

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• Rusting results from tiny galvanic cells formed by water droplets.

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Oxidation: Fe(s) → Fe²⁺(aq) + 2 e^–

Reduction: O₂(g) + 4 H⁺(aq) + 4 e^– → 2 H₂O(l)

Overall: 2 Fe(s) + O₂(g) + 4 H⁺(aq) → 2 Fe²⁺(aq) + 2 H₂O(l)

Prevention of corrosion

• Galvanizing: is the coating of iron with zinc. Zinc is more easily oxidized than iron, which protects and reverses oxidation of the iron.
• Cathodic Protection: is the protection of a metal from corrosion by connecting it to a metal (a sacrificial anode e.g. Mg or Zn) that is more easily oxidized.
• Electroplating.
• By applying paint, grease, rubber to prevent contact of metal surface from air.

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