Electrochemistry and Electrolysis – NDA Chemistry Notes

You use electrochemistry every day without knowing it. The battery in your phone, the torch that lights your way, the gold ornaments plated on cheaper metal, all of these involve electrochemistry.

Electrochemistry is the study of the relationship between chemical reactions and electricity. It covers two main areas. The first is how chemical reactions produce electricity (galvanic cells and batteries). The second is how electricity drives chemical reactions (electrolytic cells and electrolysis).

1. Electrolytes and Non-Electrolytes

Think of what happens when you dissolve salt in water. The water becomes a conductor of electricity. This is because the salt breaks into ions (Na⁺ and Cl⁻) in water, and ions carry the electric current.

An electrolyte is a substance that conducts electricity when dissolved in water or when melted. It conducts electricity because it produces ions in solution. A non-electrolyte is a substance that does NOT conduct electricity when dissolved in water. It does not produce ions in solution.

Type	Behaviour	Examples
Strong electrolyte	Fully ionises in water, excellent conductor	HCl, H₂SO₄, NaOH, NaCl
Weak electrolyte	Partially ionises, poor conductor	Acetic acid, ammonia solution
Non-electrolyte	No ionisation, does not conduct	Sugar, ethanol, urea, glucose

Pure water does NOT conduct electricity. It has very few ions.

To make pure water conduct electricity, you must add an electrolyte (something that produces ions). Lemon juice contains citric acid which ionises and conducts electricity. Kerosene, mustard oil, and sugar do NOT conduct electricity even when added to water. They do not produce ions.

2. Electrolytic Cell (Electrolysis)

Think of a torch battery connected to two metal plates dipped in saltwater. If you connect the battery, current flows. Something happens at each plate. This setup is an electrolytic cell.

An electrolytic cell uses electrical energy to drive a non-spontaneous chemical reaction. Electricity is supplied from an external source to cause a chemical reaction.

Key terms: The electrode connected to the positive terminal of the battery is the anode. Oxidation occurs at the anode. The anode loses electrons. The electrode connected to the negative terminal of the battery is the cathode. Reduction occurs at the cathode. The cathode gains electrons.

Simple memory rule: At the anode, oxidation occurs. At the cathode, reduction occurs. Another memory rule: AN OX and RED CAT. AN OX = ANode OXidation. RED CAT = REDuction at CAThode.

3. Electrolysis of Water

Pure water is a very poor conductor. But when dilute sulphuric acid or sodium hydroxide is added, water can be electrolysed.

When water is electrolysed: At the cathode (negative electrode): 2H⁺ + 2e⁻ → H₂ (hydrogen gas is produced). At the anode (positive electrode): 4OH⁻ → 2H₂O + O₂ + 4e⁻ (oxygen gas is produced).

The ratio of gases produced is H₂:O₂ = 2:1 by volume. Hydrogen collects at the cathode. Oxygen collects at the anode.

4. Electrolysis of Brine (Chlor-Alkali Process)

Brine is a concentrated solution of common salt (NaCl) in water.

When brine is electrolysed: At the cathode: 2H⁺ + 2e⁻ → H₂ (hydrogen gas). At the anode: 2Cl⁻ → Cl₂ + 2e⁻ (chlorine gas). At the cathode region: NaOH (sodium hydroxide) builds up in solution.

So the three products of electrolysis of brine are: chlorine gas, hydrogen gas, and sodium hydroxide solution. This is called the chlor-alkali process because it produces chlorine (chlor) and alkali (NaOH).

5. Electroplating

You may have seen cheap metal jewellery with a shiny gold coating. The gold coating is applied by electroplating.

Electroplating is the process of depositing a thin layer of one metal over another using electrolysis. When copper or nickel jewellery is placed in a solution containing a gold salt, a thin layer of gold is deposited on the jewellery.

How electroplating works: The object to be coated is made the cathode (connected to the negative terminal). The metal to be deposited is made the anode (connected to the positive terminal). The electrolyte is a solution of the metal salt to be deposited.

During electroplating: At the cathode, metal ions from the solution deposit on the object (reduction). At the anode, the anode metal dissolves into the solution (oxidation). The result is the object gets coated with a thin layer of the desired metal.

Uses of electroplating: Jewellery: gold and silver plating on cheaper metals. Cutlery: silver plating on steel. Car parts: chromium plating on steel for protection and shine. Electronic components: gold plating on contact points for better conductivity. Prevention of corrosion: zinc plating (galvanisation) on iron.

6. Electrolytic Refining of Metals

Electroplating principles are also used to purify metals.

In electrolytic refining: The impure metal is made the anode. Pure metal is made the cathode. The electrolyte is a solution of a salt of that metal.

When current flows, the impure anode dissolves. Pure metal deposits on the cathode. Impurities fall to the bottom as anode sludge. This is how copper is purified from impure copper ore.

7. Galvanic Cell (Voltaic Cell)

You use a galvanic cell every time you use a torch or a remote control.

A galvanic cell converts chemical energy into electrical energy. A spontaneous chemical reaction (redox reaction) inside the cell produces electricity.

In a galvanic cell: Oxidation occurs at the anode (negative electrode in a galvanic cell). Reduction occurs at the cathode (positive electrode in a galvanic cell).

Feature	Galvanic Cell	Electrolytic Cell
Energy conversion	Chemical → Electrical	Electrical → Chemical
Reaction type	Spontaneous	Non-spontaneous
Power source	Self-powered	External power supply needed
Anode polarity	Negative	Positive
Cathode polarity	Positive	Negative
Examples	Dry cell, lead-acid battery, Daniel cell	Electrolysis of water, electroplating, chlor-alkali

8. Daniel Cell

The Daniel cell was one of the earliest galvanic cells. It uses zinc and copper as electrodes.

Zinc rod (anode) in zinc sulphate (ZnSO₄) solution. Copper rod (cathode) in copper sulphate (CuSO₄) solution. A salt bridge connects the two half-cells.

At the anode: Zn → Zn²⁺ + 2e⁻ (zinc oxidises, dissolves). At the cathode: Cu²⁺ + 2e⁻ → Cu (copper deposits).

The zinc electrode slowly dissolves. Copper deposits on the copper electrode. The reaction is spontaneous. Electricity is produced.

9. Types of Cells

Primary Cell

A primary cell cannot be recharged. Once the chemicals inside are used up, the cell is discarded. Examples: dry cell (Leclanché cell), alkaline cell, mercury cell.

A dry cell produces electricity from chemical energy. The dry cell is the most common battery used in torches, remote controls, and clocks. An electrochemical cell that cannot be charged (recharged) is a primary cell. [NDA 2012-II]

Secondary Cell (Storage Cell)

A secondary cell can be recharged by passing electricity through it. During discharge (use), it acts as a galvanic cell (chemical → electrical energy). During recharge, it acts as an electrolytic cell (electrical → chemical energy). Examples: lead-acid battery, nickel-cadmium (Ni-Cd) battery, lithium-ion battery.

Lead-Acid Battery

The lead-acid battery is used in automobiles (cars, trucks, motorcycles). A typical automobile battery has six cells connected in series. Each cell has a potential (voltage) of 2 volts. Total voltage = 6 × 2 = 12 volts.

Electrodes: Lead (Pb) at the anode and Lead dioxide (PbO₂) at the cathode. Electrolyte: Dilute sulphuric acid (H₂SO₄).

During discharge (when the battery is powering the car): The battery acts as a galvanic cell. It converts chemical energy to electrical energy. During recharge (when the alternator charges the battery): The battery acts as an electrolytic cell. Electrical energy converts chemical products back to the original reactants.

Facts about the lead-acid automobile battery that are directly tested: It is a lead-acid battery. It has six cells. Each cell gives 2 V. During discharge, cells work as galvanic cells. During recharge, cells work as electrolytic cells.

Fuel Cell

A fuel cell converts chemical energy directly into electrical energy through an electrochemical reaction. It needs a continuous supply of fuel (hydrogen).

A fuel cell is NOT rechargeable. It works as long as fuel is supplied. [NDA 2012-II] [Also covered in Chapter 18]

10. Sacrificial Anode and Cathodic Protection

Ships, underground pipelines, and iron structures rust over time. One way to protect them is to attach a more reactive metal.

When a more reactive metal is attached to an iron structure, the reactive metal acts as the anode and corrodes preferentially. The iron acts as the cathode and is protected. This more reactive metal is called a sacrificial anode.

The sacrificial anode must be MORE reactive than iron (Fe) in the reactivity series. Reactivity order: Mg > Al > Zn > Fe > Sn.

Zinc, aluminium, and magnesium can all be used as sacrificial anodes. Tin (Sn) CANNOT be used as a sacrificial anode. Tin is LESS reactive than iron. If tin was used, iron would corrode instead of tin.

11. Misch Metal

Misch metal is an alloy made from rare earth metals (mainly cerium, lanthanum, and other lanthanides).

Misch metal is widely used in the manufacture of cigarette lighters. When misch metal is struck, it produces sparks. These sparks ignite the lighter fuel.

12. Faraday’s Laws of Electrolysis

Michael Faraday discovered the quantitative laws governing electrolysis. These laws connect the amount of electricity passed to the amount of substance deposited.

Faraday’s First Law: The amount of substance deposited at an electrode is directly proportional to the quantity of electricity passed through the electrolyte.

Faraday’s Second Law: When the same quantity of electricity is passed through different electrolytes, the masses of substances deposited are proportional to their equivalent weights.

At NDA level, the key understanding is: More electricity = more substance deposited. Equivalent weight determines how much of each substance is deposited per unit of electricity.

The equivalent weight of H₂SO₄ is 49 (molecular weight 98 divided by 2, because it can give 2 H⁺ ions). [NDA 2012-II] Covered in Chapter 4

13. Standard Hydrogen Electrode

The standard hydrogen electrode is the reference point for measuring all electrode potentials.

Standard electrode potential of the hydrogen electrode = 0 (by definition). This value applies at 25°C, with H₂ gas at 1 atm pressure, and H⁺ ion concentration of 1 mol/L.

The absolute electrode potential of the standard hydrogen electrode is NOT zero. The standard electrode potential (measured against itself as reference) is defined as zero.