Notes-Part-3-Class-12-Chemistry-Chapter-5-Electrochemistry-Maharashtra Board

Topics to be Learn : Part-1 Introduction Electric conduction Electrical conductance of solution Topics to be Learn : Part-2 Electrochemical cells Electrolytic cell Galvanic or voltaic cell Electrode potential and cell potential Thermodynamics of galvanic cells Topics to be Learn : Part-3 Reference electrodes Galvanic cells useful in day-to-day life Fuel cells Electrochemical series (Electromotive series)

Reference electrodes :

A single electrode potential can’t be measured but the cell potential can be measured :

Explanation :

• The electrode potential, according to Nernst theory, is the potential difference between the metal and the ionic layer surrounding it at equilibrium, i.e. the potential across the electric double layer.
• For measuring the single electrode potential, one part of the double layer, that is metallic layer can be connected to the potentiometer but not the ionic layer. Hence, single electrode potential can’t be measured experimentally.
• When an electrochemical cell is developed by combining two half cells or electrodes, they can be connected to the potentiometer and the potential difference or cell potential can be measured.

E_cell : E₂ − E₁, where E₁ and E₂ are reduction potentials of two electrodes.

• If one of the electrode potentials is known or arbitrarily assumed and E_cell is measured by potentiometer, then potential of another electrode can be obtained. Therefore it is necessary to choose a reference electrode with arbitrarily fixed potential and measure the potentials of other electrodes.
• Therefore Standard Hydrogen Electrode (SHE) is selected assuming arbitrary potential 0.0 volt. Hence potentials of all other electrodes are referred to as hydrogen scale potentials.

Standard hydrogen electrode (SHE) :

A single electrode potential cannot be measured, but the cell potential can be measured experimentally. Hence, it is necessary to have a reference electrode. S.H.E. is a primary reference electrode.

Construction :

The standard hydrogen electrode (S.H.E.) consists of a glass tube at the end of which a piece of platinised platinum foil is attached as shown in below Fig. Around this plate there is an outer jacket of glass which has a side inlet through which pure and dry hydrogen gas is bubbled at one atmosphere pressure. The inner tube is filled with a little mercury and a copper wire is dipped into it. This provides an electrical contact with the platinum foil. The outer jacket ends into a broad opening.

• The whole assembly is kept immersed in a solution containing hydrogen ions (H⁺) of unit activity.
• This electrode is arbitrarily assigned zero potential.
• The platinised platinum foil is used to provide an electrical contact for the electrode. This permits rapid establishment of the equilibrium between the hydrogen gas adsorbed by the metal and the hydrogen ions in solution.

Representation of S.H.E. :

H⁺|H₂ (g, 1atm)| Pt

(1 M)

Working :

Reduction : H⁺_(aq) + e⁻ ⇌ ½ H_2(g)  E⁰ = 0.00 V

H₂ gas in contact with H⁺_(aq) ions attains an equilibrium establishing a potential.

Applications of SHE : A reversible galvanic cell with the experimental (indicator) electrode,

Zn²⁺|Zn_(s) and SHE can be developed as follows

(1 M)

Zn |Zn²⁺|| H⁺_(aq)|H₂(g, 1 aim)|Pt

      (1 M)  (1 M)

Then in general,

E_cell = E_SHE − E_M

= 0 − E_M

= − E_M

Thus the potential can be directly obtained.

Disadvantages (Drawbacks or Difficulties) :

• It is difficult to construct and handle SHE.
• Pure and dry H₂ gas cannot be obtained.
• Pressure of H₂ gas cannot be maintained exactly at 1 atmosphere.
• The active mass or concentration of H⁺ from HCl cannot be maintained exactly unity.

Potential of hydrogen electrode:

Hydrogen gas electrode is represented as,

H⁺_(aq)|H₂ (g, 1atm)| Pt

Electrode reduction reaction is,

2 H⁺_(aq) + 2 e⁻ —> H_2(g)

By Nernst equation, the reduction potential is,

E_H2 =  \(E^0_{H_2}-\frac{0.0592}{2}log_{10}\frac{P_{H_2}}{[H^+]^2}\)

‘.’  E⁰_H2 = E⁰_SHE = 0.0 V

∴ E_H2 = \(-\frac{0.0592}{2}log_{10}\frac{P_{H_2}}{[H^+]^2}\)

If H₂ gas is passed at 1 atm, then P_H2 = 1 atm

∴ E_H2 = \(-\frac{0.0592}{2}log_{10}\frac{1}{[H^+]^2}\)

           = \(-\frac{0.0592}{2}log_{10}[H^+]^2\)

           = −0.0592 log₁₀[H⁺]

           = 0.0592[− log₁₀(H⁺)]

           = 0.0592 pH

∴ pH = E_H2/0.0592

Know This : Alkaline dry cell : The Leclanche' dry cell works under acidic conditions due to the presence of NH4Cl.The difficulty with this dry cell is that Zn anode corrodes due to its actions with H+ ions from NH4+ ions. This results in shortening the life of dry cell. To avoid this a modified or of the dry cell called alkaline dry cell has been proposed. In alkaline dry cell NaOH or KOH is used as electrolyte in place of NH4Cl. The alkaline dry cell has longer life than acidic dry cell since the Zn corrodes more slowly.

Lead storage battery (Lead accumulator) :

• Because electrical energy and emf are not generated within the cell, but are previously stored by passing an electric current, the lead accumulator is a secondary electrochemical cell. As a result, it is also known as a lead accumulator or a lead storage battery.
• It is reversible since the electrochemical reaction can be reversed by passing an electric current in opposite direction and consumed reactants can be regenerated.
• Hence battery can be charged after it is discharged.

Construction : In a lead accumulator, the negative terminal (anode) is made up of lead sheets packed with spongy lead, while the positive terminal (cathode) is made up of lead grids packed with PbO₂.

Sulphuric acid of about 38% strength (% w/w) or specific gravity 1.28 or 4.963 molar is the electrolyte in which the lead sheets and lead grids are dipped.

The positive terminal and negative terminal are alternatively arranged in the electrolyte and are separately interconnected.

Notation of the cell : The cell is formulated as Pb_(s)|PbSO_4(s)|38%H₂SO_4(aq)|PbSO_4(s)|PbO_2(s)|Pb_(s)

Working of a lead accumulator :

Discharging : When the electric current is withdrawn from lead accumulator, the following reactions take place :

Oxidation at the — ve electrode or anode :

Pb_(s) → Pb²⁺_(aq) + 2e⁻                              (oxidation)

Pb²⁺_(aq) + SO₄²⁻_(aq) → PbSO_4(s)             (precipitation)

----------------------------------------

Pb_(s) + SO₄²⁻_(aq) → PbSO_4(s) + 2e⁻  (overall oxidation)

Reduction at cathode (+) : The electrons produced at anode travel through external circuit and re-enter the cell at cathode. At cathode PbO₂ is reduced to Pb²⁺ ions in presence of H⁺ ions. Subsequently Pb²⁺ ions so formed combine with SO₄²⁻ ions from H₂SO₄ to form insoluble PbSO₄ that gets coated on the electrode.

PbO_2(s) + 4H⁺_(aq) + 2e⁻ → Pb²⁺_(aq) + 2H₂O_(l)                           (reduction)

Pb²⁺_(aq) + SO₄²⁻_(aq) → PbSO_4(s)                                          (precipitation)

--------------------------------------------------------------

PbO_2(s) + 4H⁺_(aq) + SO₄²⁻_(aq) + 2e⁻→ PbSO_4(s) + 2H₂O_(l)  (overall reduction)

Net cell reaction during discharge: The net cell reaction is the sum of overall oxidation at anode and overall reduction at cathode.

Pb_(s) + PbO_2(s) + 4H⁺_(aq) + 2SO₄²⁻_(aq)  → 2 PbSO_4(s) + 2H₂O_(l)

or

Pb_(s) + PbO_2(s) + 2H₂SO_4(aq) → 2PbSO_4(s) + 2H₂O_(l) ...(iii)

The cell potential or emf of the cell depends upon the concentration of sulphuric acid. During the operation, the acid is consumed and its concentration decreases and specific gravity decreases from 1.28 to 1.17. As a result, the emf of the cell decreases. The emf of a fully charged cell is about 2.0 V.

Recharging of the cell : When the discharged battery is connected to external electric source and a higher external potential is applied the cell reaction gets reversed generating H₂SO_4.

Reduction at the − ve electrode or cathode :

PbSO_4(s) + 2e⁻ → Pb_(s)  + SO₄²⁻_(aq)

Oxidation at the +ve electrode or anode :

PbSO_4(s) + 2H₂O_(l) → PbO_2(s) + 4H⁺_(aq) + SO₄²⁻_(aq) + 2e⁻

The net reaction during charging is

2PbSO_4(s)  + 2H₂O_(l) → Pb_(s) + PbO_2(s) + 4H⁺_(aq) + 2SO₄²⁻_(aq)

OR

2PbSO_4(s)  + 2H₂O_(l) → Pb_(s) + PbO_2(s) + 2H₂SO_4(aq)

The emf of the accumulator depends only on the concentration of H₂SO_4.

Applications :

• It is used as a source of d.c. electric supply.
• It is used in automobile in ignition circuits and lighting the head lights by connecting 6 batteries giving 12 V potential.
• It is also used in invertors.

Fuel cells :

Principle :

• The operation of fuel cells is based on the fact that combustion reactions are redox in nature and can generate electricity.
• The fuel cells differ from ordinary galvanic cells in that the reactants are not placed within the cell. They are continuously supplied to electrodes from a reservoir.
• In these cells one of the reactants is a fuel such as hydrogen gas or methanol. The other reactant such as oxygen, is oxidant.
• The simplest fuel cell is hydrogen-oxygen fuel cell.

Advantages :

• The fuel cell operates continuously as long as H₂ and O₂ gases are supplied to the electrodes.
• The cell reactions do not cause any pollution.
• The efficiency of this galvanic cell is the highest about 70 % as compared to ordinary galvanic cells.

Drawbacks of H₂-O₂ fuel cell :

• The cell requires expensive electrodes like Pt, Pd.
• In practice, voltage is less than 1.23 volt due to spontaneous reactions at the electrodes.
• H₂ gas is expensive and hazardous.

Applications of fuel cells

• The fuel cells are used on experimental basis in automobiles.
• The fuel cell are used for electrical power in the space programme.
• In space crafts the fuel cell is operated at such a high temperature that the water evaporates at the same rate as it is formed. The vapour is condensed and pure water formed is used for drinking by astronauts.
• In future, fuel cells can possibly be explored as power generators in hospitals, hotels and homes.

Electrochemical series (Electromotive series) :

Electrochemical series (Electromotive series) : It is defined as the arrangement in a series of electrodes of elements (metal or non-metal in contact with their ions) with the electrode half reactions in the decreasing order of their standard reduction potentials. .

Key points of electrochemical series :

The conventions used in the construction of electrochemical series (or electromotive series) are as follows :

• The (reduction) electrodes or half cells of the elements are written on the left hand side of the series and they are arranged in the decreasing order of their standard reduction potentials (E⁰_red).
• Reduction half reactions are written for each half cell in such a way that the species with higher oxidation state and electrons are on left hand side while reduced species with lower oxidation state are on right hand side.
• The standard reduction potential of standard hydrogen electrode is 0.00 V, i.e., E°_H2 = 0.0 V.
• The electrodes and half cell reactions with positive E⁰_red values are located above hydrogen and those with negative E⁰_red values below hydrogen. Above hydrogen, positive E⁰_red values increase, while below hydrogen negative E⁰ values increase.
• The positive E⁰_red values indicate the tendency for reduction and the negative E⁰_red values indicate the tendency for oxidation.
• The elements, whose electrodes are at the top of the series having high positive values for E⁰_red are good oxidising agents.
• The elements, whose electrodes are at the bottom of the series having high negative values for E⁰_red are good reducing agents.

General rules

• An oxidizing agent can oxidize any reducing agent that appears below it, and cannot oxidize the reducing agent appearing above it in the electrochemical series.
• An reducing agent can reduce the oxidising agent located above it in the electrochemical series.

Remember : The left side of half reaction has cations of metals or nonmetallic molecules (oxidants). There are free metals or anions of non metals on the right side (reductants). The fuel cells for power electric vehicles incorporate the proton conducting plastic membrane. These are proton exchange membranes (PEM) fuel cells.