This document provides an overview of electrochemistry concepts including:
- The differences between electrolytic and galvanic cells and how they convert between chemical and electrical energy.
- How standard electrode potentials are measured and used to determine the electrode potential of other electrodes relative to the standard hydrogen electrode.
- How galvanic cells work and the relationship between electrode potentials and the overall cell potential.
- The electrochemical series and how it arranges electrodes by their reduction potentials.
- How corrosion occurs via electrochemical reactions and how the presence of more or less reactive metals affects the corrosion of iron.
- Examples of primary batteries like dry cells and secondary batteries like lead-acid batteries, including their cell reactions.
Conductors and Non-Conductors
Substances can be classified as conductors and non-conductors based on their ability to conduct electricity.
Conductors: Substances that allow electric current to flow through them are called conductors. For example, Plastic, Wood, etc.
Non-Conductors: Non-conductors are insulators that do not allow electricity to pass through them. For example, Copper, Iron, etc.
Types of Conductors
Conductors are divided into two groups: Metallic conductors and Electrolytes.
Metallic Conductors: These conductors conduct electricity by the movement of electrons without any chemical change during the process. This type of conduction happens in solids and in the molten state.
Electrolytes: They conduct electricity by the movement of the ions in the solutions. It is present in the aqueous solution.
Distinguish between Metallic and Electrolytic Conduction
Metallic Conduction Electrolytic Conduction
The movement of electrons causes the electric current The movement of ions causes the electric current
There is no chemical reaction Ions get ionised or reduced at the electrodes
There is no transfer of matter It involves the transfer of matter in the form of ions
Follows Ohm’s law Follows Ohm’s law
Resistance increases with an increase in temperature Resistance decreases with an increase in temperature
Faraday’s law is not followed Follows Faraday’s law
Electrolytes
(a) Substances whose aqueous solutions allow the conductance of electric current and are chemically decomposed are called electrolytes.
(b) The positively charged ions furnished by the electrolyte are called cations, while the negatively charged ions furnished by the electrolyte are called anions.
Types of Electrolytes
(a) Weak electrolytes: Electrolytes that are decomposable to a very small extent in their dilute solutions are called weak electrolytes. For example, organic acids, inorganic acids and bases etc.
(b) Strong electrolytes: Electrolytes that are highly decomposable in aqueous solution and conduct electricity frequently are called electrolytes. For example, mineral acid and salts of strong acid.
Electrode
For the electric current to pass through an electrolytic conductor, the two rods or plates called electrodes are always needed. These plates are connected to the terminals of the battery to form a cell. The electrode through which the electric current flows into the electrolytic solution is called the anode, also called the positive electrode, and anions are oxidised here.
An electrode through which the electric current flows out of the electrolytic solution is called the cathode, also called the negative electrode, and cations are reduced there.
Electrolysis
Electrolysis is the process of chemical deposition of the electrolyte by passing an electric current. Electrolysis takes place in an electrolytic cell. This cell will convert the electrical energy to chemical energy.
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2. Contents :
* Electrolytic and Galvanic cell
* Galvanic cell and electrode potential
* Measurement of std. electrode potential
* Representation of half cell reaction
* Electrochemical series and its applications
* Nernst equation
* Corrosion and rusting of iron
* Commercial battery and cell.
* Exercises
3. Electrolytic cell Galvanic cell
Electrical energy converts into chemical
energy
Chemical energy converts into electrical
energy
4. Standard electrode potential
The electrode potential measured under standard condition i.e. at 25oC and
At 1 .00 atm pressure.
Standard hydrogen electrode
H++e ½ H2 , Eo = 0.00 V
Electrode potential of any electrode
may be can calculated by forming a
galvanic cell of that electrode with
hydrogen electrode.
5. E0
cell = E0
cathocde – E0
anode
E0cell = Standard electrode potential
of Galvanic cell
E0cathocde = Standard electrode
potential of cathode
E0anode = Standard electrode
potential of anode
Hydrogen electrode = cathode
‘M’ electrode = anode
E0 for hydrogen electrode = 0
E0cell = V
On applying in the formula
V = 0 – E0anode (M)
Therefore
electrode potential of ‘M’ = – v volt
6. Over all cell reaction :
Zn(s) + Cu++(aq) Zn++(aq) + Cu(s)
7. In this series standard electrode potential of various electrodes are arranged
In the increasing order of their reduction potential
8. Now it can be understood that H+ ions
reduces into H2(g) in place of Na+ ions
into Na(s)
9. Iron displaces cu from CuSO4
solution
Magnesium displaces cu from
CuSO4 solution
10.
11. Ecell = electrode potential of cell , Keq = equilibrium constant ,
∆G = gibb’s energy, T = temperature, R = gas constant ,
F = faraday’s constant
12. Chemistry of rusting of iron
Rusting is a electrochemical reaction and it become faster in the
presence of less reactive metal than iron and will not occurs in the
presence of more reactive metal than iron.
13. Primary Batteries (Dry cell)
Anode : Zn n++ + 2e-
Cathode : 2NH4
+ (aq) + 2MnO2(s) 2e- Zn++ + 2MnO(OH) + 2NH3
Overall : Zn + 2NH4+(aq) + 2MnO2(S) Zn+++ 2MnO(OH) + 2NH3
*Cell reaction can not be reversed therefore they are not chargeable.*
14. Secondary cell ( Lead storage battery)
At anode : At cathode
Pb(s) Pb++(aq) + 2e– PbO2(s) 4H+ +2e– Pb++(aq) + 2H2O
Pb++(aq) + SO4
–2(aq) PbSO4(s) Pb++(aq) + SO4
–2 (aq) PbSO4
Pb(s) +SO4
–2 (aq) PbSO4(s) PbO2(s) + 4H+ + 2e– + SO4
–2 (aq) PbSO4(s) + 2H2O(l)
Overall cell reaction :
Pb(s) + PbO2(s) + 2H2SO4(aq) 2PbSO4(s) + 2H2O (l)
*Cell reaction can be reversed therefore these cells can be recharged.*
15. Anode : 2[H2(g) + 2OH– (aq) 2H2O (l) + 2e – ]
Cathode : O2(g) + 2H2O (l) + 4e– 4OH– (aq)
Overall : 2H2(g) + O2(g) 2H2O(l)
Advantages of Fuel cell
* High efficiency
*Continuous source of energy
* Pollution free working
16. Exercises
Multiple choice questions :
Q. 1 Which of the following is not a good conductor ?
(a) Cu metal (b) Aqueous NaCl (c) Molten NaCl (d) Solid NaCl
Q.2 The reduction potential of Zn , Cu , Fe , and Ag are in the order :
(a) Ag , Cu , Fe , Zn (b) Cu , Ag , Fe , Zn (c) Zn , Cu , Fe , Ag (d) Fe , Zn , Cu , Ag.
Q.3. In an Galvanic cell which of the following statements is not correct ?
(a) Anode is negatively charged (b) Cathode is positively charged
(c) Reduction takes place at the anode (d) Reduction takes place at the cathode
Q.4. When lead storage battery discharges :
(a) SO2 is evolved (b) PbSO4 is consumed (c) Lead is formed (d) H2SO4 is consumed
Q.5. Rust is a mixture of :
(a) FeO and Fe(OH)3 (b) FeO and Fe(OH)2 (c) Fe2O3 and Fe(OH)3 (d) F3O4 and Fe(OH)3
Q.6. Galvanised iron sheets are coated with :
(a) C (b) Cu (c) Zn (d) Ni
Q.7. E0
cell and ∆G0 are related as :
(a) ∆G0 = n f E0
cell (b) ∆G = – n f E0
cell (c) ∆G0 = – n f E0
cell (d) ∆G0 = n f E0
cell = 0
By – Rajesh Trivedi ( PGT – Chemistry)
