Electrochemistry
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Unit – 5
Electrochemistry
| Oxidation | Reduction |
|---|---|
| loss of e− | gain of e− |
| +ve charge increase | −ve charge increase |
| −ve charge decrease | +ve charge decrease |
Electrochemical cell
Electrochemical cell is a device which consist of two metallic electron dipoles in same or different electrolyte. It is two types.
1) Electrolytic cell
The cell in which two metallic electrodes are placed into suitable electrolyte soln and external source of current is used to cause the chemical reaction is known as electrolytic cell.
2) Galvanic cell / Voltaic cell
Galvanic cell consist of two separate compartment each compartment consist of metallic electrode and corresponding salt soln or electrolyte. Each compartment are half cell. In left half cell oxidation take place called anode which is denoted by −ve charge. In oxidation half cell electrode lose electron. In right half cell reduction take place called cathode which is denoted by +ve charge. In reduction half cell electrode gain electron.
Salt Bridge
The salt bridge is U-shaped glass tube filled with NaCl / KNO3 / NH4NO3 gel or agar medium which maintain electrical neutrality bet’n two compartment. It also prevent direct contact bet’n oxidant and reductant but provide electric or ionic contact bet’n anode and cathode.
Representation of Galvanic cell
If concn of electrolyte is given:
Cell reaction
Oxidation half cell at anode:
At cathode:
Net rxn:
| Anode | Cathode |
|---|---|
| low | high |
| oxidation | reduction |
| loss of e− | gain of e− |
| left | right |
Single electrode potential
When metal (Zn) is dipped in metallic salt soln (ZnSO4) then an electrical double layer is formed due to accumulation of electron. One layer is formed on the surface of metal and other layer is formed on the soln. The potential difference bet’n electrode and the soln is known as single electrode potential.
Standard electrode potential
The electrode potential measures at standard condition (i.e temp = 25°C, pressure = 1 atm and [ion] = 1M) is called standard electrode potential.
EMF of cell / E cell
Evaluation of EMF of cell / E cell.
The difference bet’n electrode potential of cathode and anode is called EMF cell.
ER − EL
At standard condition:
= E°R − E°L
Standard hydrogen electrode
The electrode in which the pure hydrogen gas at 1 atm is passed to the soln containing 1M H+ ion (molarity) at 25°C is called standard hydrogen electrode.
Potential of hydrogen is taken on zero.
Application of standard hydrogen electrode (SHE)
- It is used to prepare calomel electrode, glass electrode etc.
- It is used as a reference electrode for the calculation of electrode potential of other electrode such as Cu, Zn etc.
NOTE: cell notation
Anode:
Cathode:
Cell Rxn
Anode:
Cathode:
3) Galvanic cell by using Zn – H2 electrode
Cell notation
Cell rxn:
At anode: Zn − 2e− → Zn2+
At cathode: 2H+ + 2e− → H2
Zn + 2H+ → Zn2+ + H2↑
Limitation of SHE
- It is difficult to maintain 1 atm pressure for long time.
- It is difficult to maintain 1M concn of H+ ion.
- Hydrogen electrode gets poisoned at trace amount if impurities present in it.
Electrochemical series
The arrangement of elements in order of their increasing electrode potential values is called electrochemical series. It is also called activity series.
| Electrode system | Electrode rxn | E° cell (volt) |
|---|---|---|
| Li+ / Li | Li+ + 1e− ⇌ Li | −3.05 |
| Ca2+ / Ca | Ca2+ + 2e− ⇌ Ca | −2.87 |
| Mg2+ / Mg | Mg2+ + 2e− ⇌ Mg | −2.37 |
| Al3+ / Al | Al3+ + 3e− ⇌ Al | −1.66 |
| Zn2+ / Zn | Zn2+ + 2e− ⇌ Zn | −0.76 |
| H+ / H2 | 2H+ + 2e− ⇌ H2 | 0.0 |
| Cu2+ / Cu | Cu2+ + 2e− ⇌ Cu | +0.34 |
| Ag+ / Ag | Ag+ + 1e− ⇌ Ag | +0.80 |
| F2 / F− | F2 + 2e− ⇌ 2F− | +2.87 |
Application of Electrochemical series
1) To compare the relative strength of an oxidising and reducing agent
The element having low standard reduction potential are strong reducing agent and the elements having high standard reduction potential are strong oxidising agent.
2) To identify an anode and cathode
The electrode with low standard reduction potential is anode and the electrode with high standard reduction potential is cathode.
3) To predict whether given redox rxn is feasible or not
If E° cell of rxn is +ve then the rxn is feasible and the value of standard Gibbs free energy is −ve and if E° cell of rxn is −ve then the rxn is non-feasible and the value of standard Gibbs free energy is +ve.
4) To calculate standard Gibbs free energy (ΔG°)
The standard Gibbs free energy is related to standard EMF by relation:
where,
- ΔG = standard Gibbs free energy
- n = no of mole of electron involved
- F = Faraday’s constant = 96500
- E°cell = emf of cell rxn
5) To calculate the equilibrium constant (Keq)
The Gibbs free energy is related to equilibrium constant by the relation:
= −2.303 RT log Keq
ΔG° = Standard Gibbs free energy
R = universal gas constant
T = Temp in kelvin scale
Keq = equilibrium constant.
6) To predict chemical reactivity of metals
The chemical reactivity of element (metals) decreases with the electrochemical series but an increase in the value of standard reduction potential.
7) To predict whether a metal will liberate hydrogen gas from dilute acid or not
Metals which have low reduction potential than hydrogen will displace H2 from dilute acid whereas which have high reduction potential than hydrogen can displace from dilute acid.
Cu + dil H2SO4 → No rxn
8) To predict whether a given metal will displace another metal from their salt soln or not
Metal having low reduction potential can displace other metal having high reduction potential from their salt soln.
Cu + ZnSO4 → No rxn
Commercial Batteries
If a number of cells are connected in series then the arrangement is called battery. It is of two types.
A) Primary cells [Non-Rechargeable Batteries]
A cell in which cell reaction occurs only once and cell become dead after a period of time and cell can’t be recharged is known as primary cell.
1) Dry cell [Zn-C battery]
It is a primary electrochemical cell. In this cell anode is zinc container and graphite rod act as cathode which is surrounded by powdered MnO2 and carbon.
The space bet’n electrode is filled with paste of NH4Cl and ZnCl2.
Reaction are complicated but can be written as:
Anode → Zn° → Zn2+ + 2e−
Cathode → MnO2 + NH4+ + e− → Mn2O3 + NH3 + H2O
Zn2+ + NH3 → [Zn(NH3)4]2+
- Ammonia is not liberated as gas but it combines with Zn2+ ion to form [Zn(NH3)4]2+ complex ion.
- A dry cell does not have long life because acidic NH4Cl continuous to corrode zinc container even when not in use.
- The potential difference of dry cell varies from 1.25 volt to 1.50 volt.
2) Hydrogen-Oxygen (H2-O2) Fuel Cell
The hydrogen-oxygen fuel cell is an primary electrochemical cell. The H2-O2 widely used in space buses and cars as alternative of petrol. In H2-O2 fuel cell hydrogen gives up electron at one electrode and oxygen gains electron at the other electrode. In fuel cell H2 and O2 gas are bubbled through two porous platinum coated electrode.
Electron produced at anode (−ve terminal) move round the external circuit to cathode (+ve terminal), where O2 gas gains electron to convert into hydroxide ion. The H+ ion diffuse through the membrane to the +ve electrode.
At anode:
At cathode:
Advantages of Fuel cell
- Energy produced by H2 gas in fuel cell is higher than petrol.
- Water is byproduct so that environment pollution significantly decreases.
- They are highly efficient where dropoff of energy from fuel cell to motor is direct.
B) Secondary cell [Rechargeable cell / battery]
A cell in which cell can be used again and again and can be recharge is known as secondary cell.
1) Lead storage battery / Accumulator cell
The cell contain about 40% H2SO4 is electrolyte in which metallic electrode are dipped. Lead coated with spongy lead act as anode and lead coated with lead dioxide [PbO2]. The potential difference of cell depends upon the no of electrode used. A cell containing a set of electrodes produces about emf of 2 volt. A commercial lead accumulator consist of 6 to 12 lead storage cell connected together.
Anode:
Pb(s) → Pb2+(aq) + 2e−
Pb2+ + SO42− → PbSO4↓
Net rxn:
Insoluble lead sulphate [PbSO4] formed as both anode and cathode. Water formed dilute acid and cell reaction prevented and cell should be recharged.
During recharging cell reaction takes place in opposite direction.
Cathode:
Anode:
Calomel electrode
Calomel electrode consist of a glass tube with side tube. On both side pure mercury (Hg) is placed at the bottom of tube and connected externally to the circuit with Pt-wire sealed in glass-tube. Pure mercury is covered with paste of mercury and mercurous chloride i.e. calomel (Hg2Cl2). The remaining surface of tube is filled with saturated KCl solution. Calomel electrode is connected with another electrode with the emf of calomel electrode is determined as +0.28 volt for 1M KCl at 25°C with SHE. If KCl is used as saturated it is known as saturated calomel electrode (SCE).
It is used as secondary reference electrode for determining electrode potential of unknown electrode as its potential remains constant for long time.
If calomel electrode act as cathode the cell notation will be:
If calomel electrode acts as anode the cell notation will be:
Discussion
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