Class 12 Chemistry Chapter 3 ELECTROCHEMISTRY

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NCERT Notes For Class 12 Chemistry Chapter 3 ELECTROCHEMISTRY

Class 12 Chemistry Chapter 3 ELECTROCHEMISTRY

ELECTROCHEMISTRY

It is a branch of chemistry that deals with the relationship between chemical energy and electrical energy and their inter conversions.

ELECTROCHEMICAL CELLS

• These are devices that convert chemical energy of some redox reactions to electrical energy.
• They are also called Galvanic cells or Voltaic cells. An example for Galvanic cell is Daniel cell.
• It is constructed by dipping a Zn rod in ZnSO₄ solution and a Cu rod in CuSO₄ solution.
•  The two solutions are connected externally by a metallic wire through a voltmeter and a switch and internally by a salt bridge.

A salt bridge is a U-tube containing an inert electrolyte like NaNO₃ or KNO₃ in a gelly like substance. The functions of a salt bridge are:

1. To complete the electrical circute
2. To maintain the electrical neutrality in the two half cells.

The reaction taking place in a Daniel cell is

Zn(s) + Cu²⁺(aq) → Zn²⁺ (aq) + Cu(s)

This reaction is a combination of two half reactions:

1. Cu²⁺ + 2 e^– → Cu(s) (reduc-on half reac-on)
2. Zn(s) → Zn²⁺ + 2 e^– (oxidation half reaction)

These reactions occur in two different portions of the Daniel cell. The reduction half reaction occurs on the copper electrode while the oxidation half reaction occurs on the zinc electrode. These two portions of the cell are also called half-cells or redox couples. The copper electrode may be called the reduction half cell and the zinc electrode, the oxidation half-cell.

Electrode Potential

• The tendency of a metal to lose or gain electron when it is in contact with its own solution is called electrode potential.
• When the concentrations of all the species involved in a half-cell is unity then the electrode potential is known as standard electrode potential.
• According to IUPAC convention, standard reduction potential is taken as the standard electrode potential.
• In a galvanic cell, the half-cell in which oxidation takes place is called anode and it has a negative potential.
• The other half-cell in which reduction takes place is called cathode and it has a positive potential.
  • In a cell, the electrons flow from negative electrode to positive electrode and the current flows in opposite direction.
• The potential difference between the two electrodes of a galvanic cell is called the cell potential and is measured in volts.
• The cell potential is the difference between the electrode potentials (reduction potentials) of the cathode and anode.
• The cell electromotive force (emf) of the cell is the potential difference between the two electrodes, when no current is flow through the cell.
• By convention, while representing a galvanic cell, the anode is written on the left side and the cathode on the right side.
• Metal and electrolyte solution are separated by putting a vertical line and a salt bridge is denoted by putting a double vertical line.

For Daniel cell, the cell representation is Zn(s)/Zn²⁺(aq)//Cu²⁺(aq)/Cu(s)

Under this convention the emf of the cell is positive and is given by the potential of the half-cell on the right hand side minus the potential of the half-cell on the left hand side

i.e. Ecell = Eright – Eleft Or, Ecell = ER – EL .

Measurement of Electrode Potential

• The potential of individual half-cell cannot be measured.
• We can measure only the difference between the two half-cell potentials that gives the emf of the cell.
• For this purpose a half-cell called Standard Hydrogen Electrode (SHE) or Normal Hydrogen Electrode (NHE) is used.
• It consists of a platinum electrode coated with platinum black.
• The electrode is dipped in an acidic solution of one molar concentration and pure hydrogen gas at 1 bar pressure and 298K is bubbled through it.

It is represented as Pt(s)/H₂(g)/H⁺ (aq).

By convention, the electrode potential of SHE is taken as zero.

To determine the electrode potential of an electrode, it is connected In series with the standard hydrogen electrode and the emf of the resulting cell is determined by the equation,

Ecell = ER – EL

Since the electrode potential of SHE is zero, the value of E_cell is equal to the electrode potential of the given electrode.

If the standard electrode potential of an electrode is greater than zero (i.e. +ve), then its reduced form is more stable compared to hydrogen gas. Similarly, if the standard electrode potential is negative then hydrogen gas is more stable than the reduced form of the species.

Electrochemical series

• It is a series in which various electrodes are arranged in the decreasing order of their reduction potential.
• In this table, fluorine is at the top indicating that fluorine gas (F₂) has the maximum tendency to get reduced to fluoride ions (F^–).
• Therefore fluorine gas is the strongest oxidising agent and fluoride ion is the weakest reducing agent.
• Lithium has the lowest electrode potential indicating that lithium ion is the weakest oxidising agent while lithium metal is the most powerful reducing agent in an aqueous solution.

Nernst Equation

• Nernst proposed an equation to relate the electrode potential of an electrode (or, emf of a cell) with the electrolytic concentration.

He showed that for the electrode reaction:

Since the concentration of any solid is taken as unity, the above equation becomes:

• Where E⁰ is the standard electrode potential, R is the gas constant (8.314 JK^–1 mol^–1), F is Faraday constant (96500 C mol^–1), T is temperature in Kelvin and [M ⁿ⁺] is the concentration of the species, M ⁿ⁺.
• In Daniel cell, the electrode reactions are:

Cu²⁺ + 2 e^– → Cu(s) (cathode)

Zn(s) → Zn²⁺ + 2 e^– (anode)

The electrode potentials are given as For Cathode:

On substituting the values of R (8.314 JK^–1 mol^–1), F (96500 C mol^–1) at 298K, the above equation becomes,

Equilibrium Constant from Nernst Equation

Electrochemical Cell and Gibbs Energy of the Reaction

• Electrical work done in one second is equal to electrical potential multiplied by total charge passed.
• Also the reversible work done by a galvanic cell is equal to decrease in its Gibbs energy.
• Therefore, if the emf of the cell is E and nF is the amount of charge passed, then the Gibbs energy of the reaction, ΔG = – nFE_cell
• If the concentration of all the reacting species is unity, then E_cell = E⁰_cell . So, ΔG⁰ = –nFE⁰ _cell

So, ΔG ⁰ = –nFE⁰ cell

Thus, from the measurement of E⁰_cell, we can calculate the standard Gibbs energy of the reaction.

Conductance of Electrolytic Solutions

Resistance (R): The electrical resistance is the hindrance to the flow of electrons. Its unit is ohm (Ω).

• The resistance of a conductor is directly proportional to the length of the conductor (Ɩ) and inversely proportional to the area of cross-section (A) of the conductor.

i.e. R α Ɩ /A

or, R = a constant x Ɩ/A

or, R = ρ x Ɩ/A, where ρ (rho) is a constant called resistivity.

Its unit is ohm-metre (Ω m) or ohm-centimetre (Ω cm).

1 Ω m = 100 Ω cm, 1 Ω cm = 10^-2 Ω m

Resistivity is defined as the resistance offered by a conductor having unit length and unit area of crosssection.

Conductance (G): It is the inverse of resistance.

i.e. G = 1/R.

Its unit is ohm^-1 or mho or Siemens (S)

Or, G = ƙ x A/Ɩ

Where, ƙ is called conductivity. It is defined as the conductance of a conductor having unit length and unit area of cross-section.

Its unit is ohm^-1 m^-1 or mho m^-1 or S m^-1 .

1 S cm^-1 = 100 S m^-1

1 S m^-1 = 10^-2 S cm^-1

• There are of two types of Conductance – electronic or metallic conductance and electrolytic or ionic conductance.
• Electrical conductance through metals is called metallic or electronic conductance and it is due to the movement of electrons.
• It depends on the nature and structure of the metal, the no. of valence electrons per atoms and temperature.
• For electronic conductance, when temperature increases, conduction decreases.
• The conductance of electricity by ions present in solutions is called electrolytic or ionic conductance. It depends on

1. the nature of electrolyte
2. size of the ion produced and their solvation
3. the nature of the solvent and its viscosity
4. concentration of the electrolyte and
5. temperature (As temperature increases electrolytic conduction also increases).

Note: Substances which allow the passage of electricity in molten state or in solution state are called electrolytes. On the passage of electricity, they undergo chemical decomposition.

Measurement of the conductivity of ionic solutions

We know that, conductivity G = ƙ x A/ Ɩ

So conductivity, ƙ = G x Ɩ/A

The quantity Ɩ/A is called cell constant (G*). It depends on the distance between the electrodes and their area of cross-section. Its unit is m^-1.

i.e. conductivity = conductance x cell constant

• So in order to determine the conductivity of an electrolytic solution, first determine the resistance by using a Wheatstone bridge.
• It consists of two resistances R₃ and R₄, a variable resistance R₁ and the conductivity cell having the unknown resistance R₂.
• It is connected to an AC source (an oscillator, O) and a suitable detector (a headphone or other electronic device, P). Direct current (DC) cannot be used since it causes the decomposition of the solution.

The bridge is balanced, when no current passes through the detector.

Under this condition,

Therefore, the unknown resistance,

By knowing the resistance, we get the value of conductance and conductivity.

Conductivity cell

It consists of two platinum electrodes coated with platinum black. The electrodes are separated by a distance Ɩ and their area of cross-section is A.

The cell constant of a conductivity cell is usually determined by measuring the resistance of the cell containing a solution whose conductivity is already known (e.g. KCl solution).

Molar conductivity (λm)

It is the conductivity of 1 mole of an electrolytic solution kept between two electrodes with unit area of

cross section and at a distance of unit length.

It is related to conductivity of the solution by the equation,

λm= ƙ/C (where C is the concentration of the solution)

Or, λm = 1000 ƙ/M (where M is the molarity)

The unit of molar conductivity is Ω^-1cm² mol^-1 or S cm² mol^-1.

1S m² mol^-1 = 10⁴ S cm² mol^-1

1S cm2 mol-1 = 10– 4 S m2 mol-1

Variation of conductivity and Molar conductivity with concentration (dilution)

• Both conductivity and molar conductivity change with the concentration of the electrolyte. We know that when a solution is diluted, its concentration decreases.
• For both strong and weak electrolytes, conductivity always decreases with dilution. This is because conductivity is the conductance of unit volume of electrolytic solution.
• As dilution increases, the number of ions per unit volume decreases and hence the conductivity decreases.
• For both strong and weak electrolytes, the molar conductivity increase with dilution (or decrease with increase in concentration), but due to different reasons.
• For strong electrolytes, as dilution increases, the force of attraction between the ions decreases and hence the ionic mobility increases. So, molar conductivity increases.
• When dilution reaches maximum or concentration approaches zero, the molar conductivity becomes maximum and it is called the limiting molar conductivity (λ⁰m).

For strong electrolytes, the relation between λm and concentration can be given as:

λm = λ⁰m – A√c

• Where ‘c’ is the concentration and A is a constant depends on temperature, the nature of the electrolyte and the nature of the solvent.
• All electrolytes of a particular type have the same value for ‘A’.
• For weak electrolytes, as dilution increases, the degree of dissociation increases. So the number of ions and hence the molar conductivity increases.
• The variation of λm for strong and weak electrolytes is shown in the following graphs:

• For strong electrolytes, the value of λ⁰m can be determined by the extrapolation of the graph. But for weak electrolytes, it is not possible since the graph is not a straight line.
• So their λ⁰m values are calculated by applying Kohlrausch’s law of independent migration of ions.

 

Kohlrausch’s law of independent migration of ions

• The law states that the limiting molar conductivity of an electrolyte can be represented as the sum of the individual contributions of the anion and the cation of the electrolyte.
• Thus if an electrolyte on dissociation gives n₍₊₎ cations and n_(-) anions, its limiting molar conductivity is given as:

λ0m = n(+)λ0(+) + n(-)λ0(-)

For NaCl, λ0m(NaCl) = λ0(Na+) + λ0(Cl-)

For CaCl2, λ0m(CaCl2) = λ0(Ca2+) + 2 x λ0(Cl-)

Applications of Kohlrausch’s law

1. Determination of λ⁰m of weak electrolytes

By knowing the λ⁰m values of strong electrolytes, we can calculate λ⁰m of weak electrolytes. For e.g.

we can determine the λ⁰m of acetic acid (CH₃COOH) by knowing the λ⁰m of CH₃COONa, NaCl and HCl as follows

(1) + (2) – (3) gives:

λ⁰m (CH₃COONa) + λ⁰m (HCl) – λ⁰m (NaCl) = λ⁰CH₃COO^– + λ⁰Na⁺ + λ⁰H⁺ + λ⁰Cl^– – λ⁰Na⁺ – λ⁰Cl^–

= λ⁰CH₃COOH

2) Determination of degree of dissociation of weak electrolytes

By knowing the molar conductivity at a particular concentration (λ^cm) and limiting molar conductivity (λ⁰m), we can calculate the degree of dissociation (α) as,

By using α, we can calculate the dissociation constant of acid as:

Electrolytic Cells and Electrolysis

• In an electrolytic cell, the electrical energy is converted to chemical energy.
• The dissociation of an electrolyte by the passage of electricity is called electrolysis.
• For e.g. when CuSO₄ solution is electrolysed by Cu electrodes, Cu is deposited at the cathode and Cu²⁺ ions are liberated from the anode.

Quantitative Aspects of electrolysis – Faraday’s laws

1) Faraday’s first law

• It states that the amount of substance deposited or liberated at the electrodes (m) is directly proportional to the quantity of electricity (Q) flowing through the electrolyte.

Mathematically, m α Q

          Or, m = zQ

• Where z is a constant called electrochemical equivalent (ECE).
• Z = equivalent weight/96500 But quantity of electricity is the product of current in ampere (I) and time in second (t). i.e. Q = It

Therefore, m= zIt

1 Faraday is the charge of 1 mole of electron or it is the amount of electricity required to deposit one gram equivalent of any substance.

• Its value is 96500 C/mol.
• For the deposition of 1 mole of Na, the amount of charge required = 1 F (Since Na⁺ + e^– → Na)
• For Ca, Q = 2F (since Ca²⁺ + 2e^– → Ca)

2) Faraday’s second law

• It states that when same quantity of electricity is passed through solutions of different substances, the amount of substance deposited or liberated is directly proportional to their chemical equivalence.

For e.g. when same quantity of electricity is passed through solutions of two electrolytes A and B, then

Products of electrolysis

The products of electrolysis depend on the following factors:

1. The nature of the electrolyte: The electrolyte may be in molten state or in aqueous solution state. For

e.g. if molten NaCl is electrolysed, Na is deposited at the cathode and chlorine is liberated at the anode.

NaCl → Na⁺ + Cl^–

At cathode: Na⁺ + e^– → Na At anode: Cl^– → ½ Cl₂ + e^–

If NaCl solution is electrolysed, we get H₂ gas at the cathode and Cl₂ gas at the anode.

NaCl solution contains 4 ions – Na⁺, Cl^–, H⁺ and OH^–

Cathode reaction: H⁺ + e^– → ½ H₂

Anode reaction: Cl^– → ½ Cl₂ + e^–

NaOH is formed in the solution.

(ii) The type of electrodes used: If the electrode is inert (e.g. Pt, gold, graphite etc.), it does not participate in the electrode reaction.

• While if the electrode is reactive, it also participate in the electrode reaction.

(iii) The different oxidising and reducing species present in the electrolytic cell and their standard electrode potentials.

• Some of the electrochemical processes are very slow and they do not take place at lower voltages. So some extra potential (called overpotential) has to be applied, which makes such process more difficult to occur.

For e.g. during the electrolysis of NaCl solution, the possible reactions at anode are:

Cl^– (aq) → ½ Cl₂ (g) + e^–; E⁰                            _cell = 1.36 V

2H₂O (l )→ O₂ (g) + 4H⁺(aq) + 4e^–; E⁰         _cell = 1.23 V

• At anode, the reaction with lower value of E⁰_cell is preferred and so water should get oxidised in preference to Cl^– (aq).
• However, on account of overpotential of oxygen, the first reaction is preferred and hence Cl₂ is formed at anode.

Batteries

A battery is basically a galvanic cell in which the chemical energy of a redox reaction is converted to electrical energy. They are of mainly 2 types – primary batteries and secondary batteries.

Primary cells:

These are cells which cannot be recharged or reused. Here the reaction occurs only once and after use over a period of time, they become dead E.g. Dry cell, mercury button cell etc.

1. Dry Cell

• It is a compact form of Leclanche cell. It consists of a zinc container as anode and a carbon (graphite) rod surrounded by powdered manganese dioxide (MnO₂) and carbon as cathode.
• The space between the electrodes is filled by a moist paste of ammonium chloride (NH₄Cl) and zinc chloride (ZnCl₂). The electrode reactions are:

Anode: Zn(s) → Zn²⁺ + 2e^–

Cathode: MnO₂+ NH₄⁺+ e^–→ MnO(OH) + NH₃

Ammonia produced in this reaction forms a complex with Zn²⁺ and thus corrodes the cell. The cell has a potential of nearly 1.5 V.

2. Mercury cell

• Here the anode is zinc – mercury amalgam and cathode is a paste of HgO and carbon.
• The electrolyte is a paste of KOH and ZnO. The electrode reactions are:

Anode reaction: Zn(Hg) + 2OH^– → ZnO(s) + H₂O + 2e^–

Cathode reaction: HgO + H₂O + 2e^– → Hg(l ) + 2OH^– The overall reaction is : Zn(Hg) + HgO(s) → ZnO(s) + Hg(l )

• The cell has a constant potential of 1.35 V, since the overall reaction does not involve any ion in solution.

b) Secondary cells

• A secondary cell can be recharged and reused again and again. Here the cell reaction can be reversed by passing current through it in the opposite direction.
• The most important secondary cell is lead storage cell, which is used in automobiles and invertors.
• It consists of lead as anode and a grid of lead packed with lead dioxide (PbO₂) as the cathode. The electrolyte is 38% H₂SO₄ solution.

The cell reactions are:

Anode: Pb(s) + SO₄^2- (aq) → PbSO₄(s) + 2e-

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

The overall cell reaction is: Pb(s)+PbO₂(s)+2H₂SO₄(aq) → 2PbSO₄(s) + 2H₂O(l)

On charging the battery, the reaction is reversed and PbSO₄(s) on anode and cathode is converted into Pb and PbO₂, respectively.

Another example for a secondary cell is nickel – cadmium cell. Here the overall cell reaction is Cd (s)+2Ni(OH)₃ (s) → CdO (s) +2Ni(OH)₂ (s) +H₂O(l )

Fuel cells

• These are galvanic cells which convert the energy of combustion of fuels like hydrogen, methane, methanol, etc. directly into electrical energy.
• One example for fuel cell is Hydrogen – Oxygen fuel cell, which is used in the Apollo space programme.
• Here hydrogen and oxygen are bubbled through porous carbon electrodes into concentrated aqueous sodium hydroxide solution.
• To increase the rate of electrode reactions, catalysts like finely divided platinum or palladium metal are filled into the electrodes.

The electrode reactions are:

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

Anode: 2H₂ (g) + 4OH^–(aq) → 4H₂O(l) + 4e^– Overall reaction is: 2H₂(g) + O₂(g) ⎯→ 2 H₂O(l )

Advantages of Fuel cells

1. The cell works continuously as long as the reactants are supplied.
2. It has higher efficiency as compared to other conventional cells.
3. It is eco-friendly (i.e. pollution free) since water is the only product formed.
4. Water obtained from H₂ – O₂ fuel cell can be used for drinking.

Corrosion

• It is the process of formation of oxide or other compounds of a metal on its surface by the action of air, water-vapour, CO₂ etc.
• Some common examples are: The rusting of iron, tarnishing of silver, formation of green coating on copper and bronze (verdigris) etc.
• Most familiar example for corrosion is rusting of iron. It occurs in presence of water and air.
• It is a redox reaction.
• At a particular spot of the metal, oxidation takes place and that spot behaves as anode. Here Fe is oxidized to Fe²⁺.

2 Fe (s)→2 Fe²⁺ + 4 e^–

• Electrons released at anodic spot move through the metal and go to another spot on the metal and reduce oxygen in presence of H⁺ .
• This spot behaves as cathode. The reaction taking place at this spot is: O₂(g) + 4 H⁺(aq) + 4 e^– ⎯→ 2 H₂O (l )
• The overall reaction is: 2Fe(s)+O₂(g) + 4H⁺(aq) → 2Fe^{2 +(}aq)+ 2 H₂O (l )
• The ferrous ions (Fe²⁺) are further oxidised to ferric ions (Fe³⁺) and finally to hydrated ferric oxide (Fe₂O₃. x H₂O), which is called rust.

Methods to prevent corrosion

1. By coating the metal surface with paint, varnish etc.
2. By coating the metal surface with another electropositive metal like zinc, magnesium etc. The coating of metal with zinc is called galvanisation and the resulting iron is called galvanized iron.
3. By coating with anti-rust solution.
4. An electrochemical method is to provide a sacrificial electrode of another metal (like Mg, Zn, etc.) which corrodes itself but saves the object (sacrificial protection).