- Alloys are homogeneous mixtures of metals (or metals with non-metals) with different properties than pure metals—often harder, less conductive, and with lower melting points.
- 24-carat gold is soft, so 22-carat gold (an alloy with copper or silver) is used in jewellery to improve strength.
- Solder, an alloy with a low melting point, is used to join electrical wires.
- Corrosion prevention methods include painting, galvanising, anodising, electroplating, and alloying.
- Galvanising, tinning, anodisation, and electroplating all involve protective coatings to prevent corrosion, while alloying enhances durability and resistance.
Definitions [20]
Definition: Electrochemistry
Electrochemistry is the study of the production of electricity from energy which is released during spontaneous chemical reactions, as well as the use of electrical energy to bring about non-spontaneous chemical transformations.
Definition: Redox Reactions
Any reaction that involves both oxidation and reduction occurring simultaneously is called an oxidation-reduction reaction or simply a redox reaction.
Define cathode
The electrode at which the reduction occur is called cathode.
Define anode
The electrode at which the oxidation occur is called anode.
Define the following term:
Fuel cell
Fuel cells are the galvanic cells in which the energy of combustion of fuels like hydrogen, methanol, etc., is directly converted into electrical energy.
Define the following term:
anion
Atoms that carry negative charge are called anions.
Define the following term:
Cation
Atoms that carry positive charge are called cations.
Definition: Electrode Potential
The potential difference developed between the electrode and the electrolyte due to the loss or gain of electrons by the electrode is called the electrode potential.
Oxidation potential: M(s) ⇌ Mⁿ⁺(aq) + ne⁻ Reduction potential: Mⁿ⁺(aq) + ne⁻ ⇌ M(s) Reduction potential = − Oxidation potential
Definition: Standard cell potential (E°cell)
It is the cell potential when the concentrations of all the species are 1 M at 25°C, and the pressure of the gas involved is 1 atm at 25°C.
Definition: Standard electrode potential (E°)
Standard electrode potential is the potential associated with the electrode reaction at an electrode when all solutes are 1 M, and all gases are at 1 atm and at 25°C (298 K).
Define standard electrode potential.
Standard electrode potential is the difference of electrical potential between a metal electrode and the solution around it at equilibrium when all the substances involved are in their standard states.
The potential of an electrode assembly is referred to as the standard electrode potential when the following conditions are satisfied.
- The temperature of the electrode assembly is 298 K (25°C).
- The ion solution used in the assembly is of concentration 1 mol L−1.
- The pressure of the gas, if used in the assembly, is 1 atm.
Definition: Cell Potential
The cell potential of the cell is the algebraic sum of the electrode potentials. Cell potential is called the electromotive force (e.m.f.) of the cell when no current is drawn through the cell.
Definition: Electrochemical Series
Based on the ease with which atoms of metals lose electrons to form positively charged ions, the metals are arranged in a series known as the electrochemical series.
or
The arrangement of electrodes in order of their decreasing standard reduction potentials is called electrochemical series.
Define electrochemical series.
The standard potentials of a number of electrodes have been determined using standard hydrogen electrodes. These electrodes with their half reactions are arranged according to their decreasing standard potentials; this arrangement is called an electrochemical series.
Define cell constant.
Cell constant is the ratio of the distance between the electrodes divided by the area of cross-section of the electrode. It is denoted by b.
Thus, Cell constant = b =`l/a`. It is expressed in unit m−1.
Define conductivity for the solution of an electrolyte.
It is the inverse of resistance R and may be simply defined as the speed through which current flows in a conductor.
c = `1/R = A/(pl)`
k = `A/l`
Here k is the specific conductance. The SI unit of conductance is Siemens, which is denoted by the symbol ‘S’ and is equal to ohm−1 or Ω−1.
Define “Molar conductivity”.
Molar conductivity is the conductance of a volume of solution containing 1 mole of dissolved electrolyte when placed between two parallel electrodes 1 cm apart and large enough to contain between them all the solution.
The conductivity, which is shown by all the ions when 1 mol of electrolyte is dissolved in the solution, is called molar conductivity; it is expressed by ∧m (lambda). If 1 mol of electrolyte is present in Vm cm3 of electrolyte solution, then ∧m = κ × V
= `(kappa xx 1000)/"Molarity" = (kappa xx 1000)/M`
Its unit is ohm−1 cm2 mol−1 or S cm2 mol−1.
Define limiting molar conductivity.
The limiting molar conductivity of an electrolyte is defined as its molar conductivity when the concentration of the electrolyte in the solution approaches zero.
When the concentration of an electrolytic solution placed between electrodes of a conductivity cell placed at a unit distance having an area of cross-section sufficient to accommodate enough volume of solution containing one mole of electrolyte approaches zero, then the conductance of the solution is known as limiting molar conductivity.
Definition: Fuel Cell
A fuel cell is a galvanic cell in which the reactants are not placed within the cell, but are continuously supplied from outside, where one reactant acts as a fuel (such as hydrogen or methanol) and the other as an oxidant (such as oxygen).
Definition: Corrosion
Corrosion is the gradual damage of metals caused by their reaction with components of the atmosphere, such as oxygen and moisture.
Formulae [1]
Write the Nernst equation and explain the terms involved.
Nernst equation can be given as,
`E = E^circ - (2.303 RT)/(nF) log_10 [["Products"]]/[["Reactants"]]`
where,
E° = Standard potential of electrode or cell,
n = Number of moles of electrons used in reaction,
F = Faraday = 96500 C/mol e−,
[Products] = Concentration of products,
[Reactants] = Concentration of reactants,
T = Temperature in K and
R = Gas constant = 8.314 J K−1 mol−1
Theorems and Laws [2]
State Kohlrausch Law.
Kohlrausch law states that at infinite dilution of the solution, each ion of electrolyte migrates independently of its co-ions and contribute independently to the total molar conductivity irrespective of the nature of other ion.
State Kohlrausch’s law of independent migration of ions.
Kohlrausch’s law states that the molar conductivity of an electrolyte at infinite dilution is the same as the sum of the anions' and cations' limited molar conductivities.
`∧_m^° = v_+ λ_+^° + v_- λ_-^°`
Here `λ_+^°` and `λ_-^°` are limiting molar conductivities of cations and anions.
Key Points
Key Points: Redox Reactions
Redox Reactions:
- A substance that oxidises another substance (and is itself reduced) is called an oxidising agent.
- A substance that reduces another substance (and is itself oxidised) is called a reducing agent.
What is Oxidation and Reduction?
| Perspective | Oxidation | Reduction |
|---|---|---|
| In terms of oxygen | Gain of one or more O atoms | Loss of one or more O atoms |
| In terms of hydrogen | Loss of hydrogen | Gain of hydrogen |
| In terms of electropositive element | Loss of electropositive element | Gain of electropositive element |
| In terms of electronegative element | Gain of electronegative element | Loss of electronegative element |
| In terms of electrons | Loss of electrons | Gain of electrons |
| In terms of oxidation number | Increase in oxidation number | Decrease in oxidation number |
Redox in Terms of Electron Transfer:
A reaction in which electrons are lost by one substance and gained by another is called a redox reaction.
- Oxidising agent = electron acceptor
- Reducing agent = electron donor
Example:
\[\mathrm{Hg}_2^{2+}+\mathrm{Sn}^{2+}\to\mathrm{Hg}+\mathrm{Sn}^{4+}\]
(Hg₂²⁺ gains electrons → reduced; Sn²⁺ loses electrons → oxidised)
Key Points: Electrochemical Cells
| Type | Electrolytic Cell | Galvanic (Voltaic) Cell |
|---|---|---|
| Energy conversion | Electrical → Chemical | Chemical → Electrical |
| Nature of reaction | Non-spontaneous | Spontaneous |
| Anode | Positive | Negative |
| Cathode | Negative | Positive |
| Electron flow | Cathode → Anode | Anode → Cathode |
| Salt bridge | Not required | Required |
Electrolysis of NaCl
1. Molten NaCl:
-
Oxidation: Cl⁻ → Cl₂ (gas)
-
Reduction: Na⁺ → Na (metal)
-
Products: Na (cathode), Cl₂ (anode)
2. Aqueous NaCl:
-
Oxidation: Cl⁻ → Cl₂
-
Reduction: H₂O → H₂ + OH⁻
-
Products: H₂ (cathode), Cl₂ (anode), NaOH formed
Key Points: Galvanic or Voltaic Cell
Components of a Galvanic Cell
| Component | Key Points |
|---|---|
| Electrodes | Surfaces where oxidation and reduction occur may be inert or active |
| Anode | Electrode where oxidation occurs; in a galvanic cell → negative electrode |
| Cathode | Electrode where reduction occurs; in a galvanic cell → positive electrode |
| Electrolyte | Substance that ionises in solution or molten state; provides ions for conduction; placed in separate containers (half-cells) |
| Salt Bridge (Structure) | U-shaped tube with electrolyte |
| Salt Bridge (Functions) | Completes electrical circuit; maintains electrical neutrality; prevents mixing of solutions |
6. Cell Notation
-
Anode written on the left, cathode on the right
-
Example:
Cu(s) | Cu²⁺(aq) || Ag⁺(aq) | Ag(s)
-
Single line (|) → phase boundary
-
Double line (||) → salt bridge
Key Points: Prediction of Reaction
| ΔG° | E°cell | Nature |
|---|---|---|
| ΔG° < 0 | E°cell > 0 | Spontaneous |
| ΔG° = 0 | E°cell = 0 | Equilibrium |
| ΔG° > 0 | E°cell < 0 | Non-spontaneous |
Key Points: Fuel Cells
Reactions
- Anode:
2H₂ + 4OH⁻ → 4H₂O + 4e⁻ - Cathode:
O₂ + 4H₂O + 4e⁻ → 4OH⁻ - Overall reaction:
2H₂ + O₂ → 2H₂O
Applications
- Spacecraft (electric power)
- Power generators (homes, hospitals)
- Automobiles (experimental)
- Clean energy for industries
Drawbacks
- Hydrogen gas is hazardous
- High cost of hydrogen preparation
Key points: Prevention of Corrosion
Concepts [25]
- Introduction to Electrochemistry
- Concept of Redox Reactions
- Electrochemical Cells
- Electrodes
- Galvanic or Voltaic Cell
- Electrode Potential and Cell Potential
- Galvanic Cells - Measurement of Electrode Potential
- Electrochemical Series (Electromotive Series)
- Relation Between Gibbs Energy Change and Emf of a Cell
- Nernst Equation - Introduction
- Nernst Equation
- Equilibrium Constant from Nernst Equation
- Electrochemical Cell and Gibbs Energy of the Reaction
- Conductance of Electrolytic Solutions - Introduction
- Conductance of Electrolytic Solutions
- Measurement of the Conductivity of Ionic Solutions
- Variation of Conductivity and Molar Conductivity with Concentration
- Electrolytic Cells and Electrolysis - Introduction
- Products of Electrolysis
- Batteries
- Primary Batteries
- Secondary Batteries
- Fuel Cells
- Corrosion of Metals
- Prevention of Corrosion
