5. Electrochemistry

Have you ever wondered how your smartphone battery stores energy, or how copper metal is purified to make electrical wires? The answer lies in a fascinating branch of chemistry called electrochemistry.

In simpler terms, electrochemistry studies how chemical energy and electrical energy are interconverted. It is at the heart of our modern world — powering our phones, refining metals, and even providing clean energy through fuel cells.

Applications of Electrochemistry

• Purification of precious metals like Cu, Ag, etc.
• Manufacture of Ca, Na, Al, etc.
• Manufacture of NaOH, Cl₂, F₂, etc.
• Fuel cells and batteries.

Before we dive into electrochemical cells, let us first understand how electricity flows through substances.

Conductors	Non-conductors (Insulators)
Substances that allow the passage of electricity.	Substances that do not allow the passage of electricity.
Examples: Metals (Cu, Fe, Ag), acids (HCl, H₂SO₄), alkalis (NaOH, KOH), salt solutions (NaCl, MgSO₄).	Examples: Sugar, glucose, phosphorus (P), sulfur (S), rubber, wood.

Conductors are classified into two types based on how they conduct electricity:

Metallic Conductors	Electrolytic Conductors
Conduct electricity through the movement (migration) of free electrons.	Conduct electricity through the movement (migration) of ions.
No chemical change occurs during conduction.	Chemical changes occur at the electrode surfaces during conduction.
Examples: Metals (Cu, Ag, Au), alloys, graphite.	Examples: Solutions of acids, bases, salts; molten salts.

Electrochemical Cells

Or, simply, electrochemical cells are devices for the interconversion of electrical and chemical energy. There are two main types:

Electrolytic Cell	Galvanic (Voltaic) Cell
Converts electrical energy → chemical energy.	Converts chemical energy → electrical energy.
Uses an external source of electricity (DC).	The spontaneous chemical reaction generates electricity.
Reaction is non-spontaneous.	Reaction is spontaneous (redox reaction).
Example: Electrolysis of NaCl.	Example: Daniel’s cell (Zn-Cu cell).

Electrolytic Cell (Voltameter)

The electrochemical cell in which the electrical energy is used to bring about a chemical reaction.

Working:

• Two metallic electrodes are dipped into a suitable electrolytic solution.
• The electrodes are connected to an external DC power source (battery).
• The electrode connected to the positive terminal of the battery is called the anode (where oxidation occurs).
• The electrode connected to the negative terminal of the battery is called the cathode (where reduction occurs).
• When electricity is passed, cations move towards the cathode and anions move towards the anode, causing chemical reactions.

Example: electrolysis of molten NaCl.

Reactions involved:

NaCl → Na⁺ + Cl⁻ (Molten)

At anode: Cl⁻ − 1e⁻ → Cl (oxidation)   Cl + Cl → Cl₂

At cathode: Na⁺ + 1e⁻ → Na (reduction)

The electrochemical cell in which electrical energy is produced from spontaneous chemical (redox) reactions.

In a voltaic cell, oxidation and reduction reactions occur in separate compartments called half-cells.

▶ Related video: youtu.be/qpFC_Ecu_yQ

Daniel’s Cell — A Classic Example (also called Zn-Cu cell)

Arrangement:

• A Zn electrode is dipped into a 1.0 M ZnSO₄ solution in one beaker.
• A Cu electrode is dipped into a 1.0 M CuSO₄ solution in another beaker.
• The electrodes are connected externally through a wire (and an ammeter/voltmeter).
• The two solutions are connected internally by the salt bridge.

Working:

• Zn gets oxidised (loses 2 e⁻ per atom) into Zn²⁺ and goes into the solution.
• These electrons travel through the external circuit to the other electrode (Cu).
• At the cathode, those electrons are accepted by Cu²⁺ ions (from CuSO₄), which are reduced into Cu metal.
• The Zn electrode is known as the anode or oxidation half. It is a −ve terminal.
• Since reduction occurs at the Cu electrode, it is called the cathode or reduction half. It is the positive terminal.
• Cations (K⁺) from the salt bridge move into the CuSO₄ beaker and anions (NO₃⁻) move into the ZnSO₄ beaker to maintain charge balance.

The reactions occurring are:

Zn → Zn²⁺ + 2e⁻  (anodic / oxidation half-reaction)

Cu²⁺ + 2e⁻ → Cu  (cathodic / reduction half-reaction)

Zn + Cu²⁺ → Zn²⁺ + Cu  (Net cell reaction)

Cell Notation

Rules for writing cell notation:

A galvanic cell is represented in a shorthand notation called cell notation. The rules are:

1. A single vertical line (|) separates the metal electrode and ions in the electrolytic solution, e.g., Zn|Zn²⁺(aq)    Cu²⁺(aq)|Cu(s) — oxidation half-cell, reduction half-cell.
2. The anode half-cell is always written on the left, and the cathode half-cell on the right.
3. A double vertical line (||) represents the salt bridge, written in the middle of two half-cells. E.g., Zn(s)|Zn²⁺(aq) || Cu²⁺(aq)|Cu(s)
4. The concentration of the electrolytic solution is written inside brackets. E.g., Zn²⁺(1M), Cu²⁺(0.5M)

When a metal (electrode) (M) is immersed in a solution containing its own ions (e.g., Zn rod in ZnSO₄ solution), the following reactions occur:

Oxidation: some metal atoms lose electrons and enter the solution as positive ions.

Reduction: Some metal ions in the solution gain electrons and deposit on the metal surface.

M ⇌ Mⁿ⁺(aq.) + n e⁻

These two processes continue until a dynamic equilibrium is established. But these two processes do not occur at the same rate initially. As a result, a layer of charge develops at the surface of the metal. This attracts oppositely charged ions from the solution, forming an electrical double layer at the interface of the metal and electrolyte solution.

Due to this separation of charge across the metal–solution interface, a potential difference is produced between the electrode and the electrolyte. This potential difference is called the electrode potential or single electrode potential.

The magnitude of electrode potential depends on the nature of the electrode, the concentration of its ions, and the temperature of the system.

Types of Electrode Potential

Oxidation Potential	Reduction Potential
Potential is developed when the electrode loses electrons (oxidation).	Potential is developed when the electrode gains electrons (reduction).
Example: Zn → Zn²⁺ + 2e⁻	Example: Cu²⁺ + 2e⁻ → Cu

Furthermore, the reduction potential of Cu²⁺/Cu = +0.34 V means that when a standard Cu electrode is connected to the standard hydrogen electrode, Cu²⁺ gets reduced to Cu, delivering 0.34 V electric potential in the cell.

Electrode System	Standard Electrode (Reduction) Potential (E°)	Tendency
Cu²⁺/Cu	+0.34 V	More tendency to gain electrons compared to SHE
Zn²⁺/Zn	−0.76 V	Less tendency to gain electrons compared to SHE

Similarly, the oxidation potential is measured at 1 M concentration, 1 atm pressure, and 298K temperature, and is called the standard oxidation potential.

The single electrode (standard) potentials of all electrodes are expressed with respect to the Standard Hydrogen Electrode (SHE), whose potential is arbitrarily taken to be equal to 0 volts at all temperatures.

The electrode tending to lose electrons more than the standard hydrogen electrode (taken as 0 V) will have a positive value of oxidation potential, and the electrode having a tendency to gain electrons greater than that of the SHE will have a positive value of reduction potential.

However, the electrode potential is usually expressed in standard reduction potential.

Factors Affecting the Magnitude of Electrode Potential

1. Nature of metal/electrode: Highly reactive metals (K, Na, Ca, etc.) have a stronger tendency to lose electrons, hence have a higher (more negative) electrode potential. Less reactive metals (Cu, Ag, Au, etc.) have less tendency to lose electrons, hence have a lower (more positive) electrode potential.
2. The concentration of metal ions: Electrode potential depends on ion concentration. The exact effect depends on the electrode reaction.
3. Temperature: Electrode potential changes with temperature, but the effect depends on the particular electrode system.

To measure the electrode potential of any given electrode, a cell should be constructed with two electrodes. One of them will act as a cathode and another as an anode.

E_cell = E_RP(Cathode) − E_RP(Anode)

To know the electrode potential of one, we must know the electrode potential of the other. Electrodes through which a cell is constructed may be indicator electrodes or reference electrodes.

The indicator electrode: the one whose potential is to be determined.

The reference electrode: one with a known potential.

Reference electrodes could be Primary reference electrodes, like SHE, and secondary reference electrodes, like Calomel electrodes.

Primary Reference Electrode	Secondary Reference Electrode
Potential is defined by international convention.	Calibrated against the primary reference electrode.
Example: Standard Hydrogen Electrode (SHE).	Example: Calomel Electrode.
E° = 0.000 V (by definition).	Has a known, constant potential.

Standard Hydrogen Electrode (SHE)

The Standard Hydrogen Electrode (SHE) is the primary reference electrode. Its potential is assigned a value of exactly 0.000 V at all temperatures, by international convention.

Construction:

• A platinum (Pt) wire is sealed inside a glass tube, with a Pt foil (coated with finely divided platinum black) attached at the bottom.
• This Pt electrode is immersed in a 1 M HCl solution (which provides H⁺ ions at 1 M concentration).
• Pure H₂ gas at 1 atm pressure is continuously bubbled over the Pt electrode surface at 298 K.
• The Pt surface acts as a reaction site for the hydrogen half-reaction.

The following reactions occur in this half of the cell, depending on whether SHE acts as an anode or a cathode.

If SHE acts as anode:  H₂(g) → 2 H⁺(aq) + 2 e⁻  (Oxidation half-reaction)

If SHE acts as the cathode:  2 H⁺(aq) + 2 e⁻ → H₂(g)  (Reduction half-reaction)

Short notation for SHE:

Pt/H₂(g)/H⁺(aq)    (1 atm) (1 M)

The standard hydrogen electrode is used to determine the electrode potential of any electrode system by connecting it with SHE, measuring the EMF of the cell, and using the preceding formula. The value of SHE is arbitrarily taken to be 0.

Disadvantages of SHE

• It is difficult to obtain pure H₂ gas.
• Maintaining a constant 1 atm pressure of H₂ throughout the experiment is challenging.
• The Pt electrode can get poisoned (contaminated) by impurities, changing its potential.
• The whole arrangement is bulky and difficult to transport.

Calomel Electrode

Due to the practical difficulties of SHE, a secondary standard electrode — the calomel electrode — is widely used in laboratories.

A platinum wire sealed in a glass tube makes contact with a layer of liquid mercury. This mercury is covered by a paste of mercury and calomel (Hg₂Cl₂). The rest of the tube is filled with a KCl solution (1N or saturated).

The reaction involved is:

Hg₂Cl₂ + 2e⁻ ⇌ 2Hg + 2Cl⁻

The electrode process is reversible, and the value of potential depends on the concentration of the KCl solution.

Type	E°
Saturated calomel electrode (saturated KCl solution)	0.2415 V
1 N calomel electrode (1 N KCl solution)	0.2800 V
0.1 N calomel electrode (0.1 N KCl solution)	0.3338 V

Advantages of Calomel Electrode

• Easy to set up and attains equilibrium quickly.
• Compact in size — easy to transport and handle.
• Can be easily connected to the salt bridge.
• Reversible — can function as both anode and cathode.
• Delivers constant, stable potential throughout the process.

Element / Couple	Equilibrium (Oxidants ⇌ Reductants)	E° (volts)
Lithium	Li⁺(aq) + e⁻ ⇌ Li(s)	−3.03
Potassium	K⁺(aq) + e⁻ ⇌ K(s)	−2.92
Calcium	Ca²⁺(aq) + 2e⁻ ⇌ Ca(s)	−2.87
Sodium	Na⁺(aq) + e⁻ ⇌ Na(s)	−2.71
Magnesium	Mg²⁺(aq) + 2e⁻ ⇌ Mg(s)	−2.37
Aluminum	Al³⁺(aq) + 3e⁻ ⇌ Al(s)	−1.66
Manganese	Mn²⁺(aq) + 2e⁻ ⇌ Mn(s)	−1.18
Zinc	Zn²⁺(aq) + 2e⁻ ⇌ Zn(s)	−0.76
Iron	Fe²⁺(aq) + 2e⁻ ⇌ Fe(s)	−0.44
Nickel	Ni²⁺(aq) + 2e⁻ ⇌ Ni(s)	−0.25
Tin	Sn²⁺(aq) + 2e⁻ ⇌ Sn(s)	−0.14
Lead	Pb²⁺(aq) + 2e⁻ ⇌ Pb(s)	−0.13
Hydrogen	2H⁺(aq) + 2e⁻ ⇌ H₂(g)	0.00
Copper	Cu²⁺(aq) + 2e⁻ ⇌ Cu(s)	+0.34
Iron(III)/Iron(II)	Fe³⁺(aq) + e⁻ ⇌ Fe²⁺(aq)	+0.77
Silver	Ag⁺(aq) + e⁻ ⇌ Ag(s)	+0.80
Mercury	Hg²⁺(aq) + 2e⁻ ⇌ Hg(s)	+0.85
Gold	Au³⁺(aq) + 3e⁻ ⇌ Au(s)	+1.50

Table: A portion of the electrochemical series — metal reducing activity increases going up; metal oxidising (electron-gaining) activity increases going down.

Applications of the Electrochemical Series

1. Predicting Relative Strengths of Oxidising and Reducing Agents

In the electrochemical series, the substances are arranged in increasing order of reduction potentials (i.e., increase in oxidising power).

• The elements at the top have a minimum tendency to get reduced (less powerful oxidising agents).
• The elements at the bottom have the maximum tendency to get reduced (more powerful oxidising agents).
• On moving from top to bottom, the tendency to get reduced increases (tendency as the oxidising agent increases).
• The more positive the standard reduction potential for an element, the more easily it gets reduced and acts as a strong oxidising agent.

2. Predicting Displacement Reactions

An electropositive metal (higher in the electrochemical series, more negative E°), displaces the less electropositive metal from the solution, as free metal precipitate.

Fe(s) + CuSO₄(aq) → FeSO₄(aq) + Cu(s)

The metal present on top of another metal in the electrochemical series is more electropositive than the other metal. (Fe is present above Cu in the electrochemical series).

The metal having a more negative value of reduction potential is more electropositive.

Similarly, the more electronegative non-metal displaces the less electronegative non-metal ion from its salt solution.

3. Calculating the EMF of the Cell

EMF of a cell is the difference between the reduction potential of the cathode and the anode.

E_cell = E_(Cathode) − E_(Anode)

4. Predicting the Feasibility of a Redox Reaction

A galvanic cell is constructed from the reaction, considering it as the cell reaction. By using the formula E°_cell = E°_cathode − E°_anode, the potential for a galvanic cell is determined.

Condition	Interpretation
E°cell = +ve	The reaction is feasible in the forward direction
E°cell = −ve	The reaction is not feasible in the forward direction
E°cell = 0	The reaction is at equilibrium

Zn-Cu Cell

• Daniel’s cell is constructed by the combination of a Zn electrode and a Cu electrode.
• It is given that the standard reduction potentials of the electrodes are E°_Zn²⁺/Zn = −0.76V and E°_Cu²⁺/Cu = +0.34V.
• The standard reduction potential of the Cu electrode is higher than that of the Zn electrode. Therefore, the Cu electrode acts as a cathode, and the Zn electrode acts as an anode.
• A Cu electrode is prepared by dipping a Cu rod in a 1M CuSO₄ solution in a beaker, and a Zn electrode is prepared by dipping a Zn rod in a 1M ZnSO₄ solution in another beaker.
• The cell is then constructed by connecting the electrodes internally by means of a salt bridge and externally by means of a wire to a voltmeter.

Cell notation:

Zn(s) | Zn²⁺(1M) || Cu²⁺(1M) | Cu(s)

Anode        salt bridge        Cathode

Reactions involved:

At Zn-anode (−ve pole): Zn − 2e⁻ → Zn²⁺  (oxidation half)

At Cu-cathode (+ve pole): Cu²⁺ + 2e⁻ → Cu  (reduction half)

Overall reaction: Zn + Cu²⁺ → Zn²⁺ + Cu

This spontaneous cell reaction brings a flow of electrons from anode to cathode across the external circuit, as a result of which an electric force develops called electromotive force (emf) or cell potential.

Standard emf/cell potential of Zn-Cu/Daniel cell:

E°_cell = E°_cathode − E°_anode = 0.34V − (−0.76V) = 0.34V + 0.76V = 1.1 V

Ag-Cu Cell

• It is known that the standard reduction potential of Ag⁺/Ag = +0.80V and that of Cu²⁺/Cu = +0.34V.
• Since the standard reduction potential of the Ag-electrode is higher than that of the Cu-electrode, the Ag-electrode acts as a cathode and Cu-electrode acts as an anode.
• An Ag electrode is prepared by dipping an Ag rod in a 1M AgNO₃ solution, and a Cu-electrode is prepared by dipping a Cu rod in a 1M CuSO₄ solution.
• The cell is then constructed by connecting the electrodes internally by means of a salt bridge and externally by means of a wire to a voltmeter.

The difference between the electrode potential of two half-cells (oxidation half and reduction half) is known as the emf of the cell or cell potential. It is calculated as:

E_cell = E_RP(Cathode) − E_RP(Anode)

During the calculation of the emf of a galvanic cell, the reduction potential is used. If oxidation potential is given, its sign should be changed to get the reduction potential.

Relationship Between Cell Potential and Free Energy

In electrochemical cells, the chemical energy (obtained through spontaneous redox reactions) is converted into electrical energy. The cell potential (E) is related to the Gibbs free energy change. The system works by transferring electrical energy through an electric circuit.

Electrical work done is equal to the electric potential (E_cell) multiplied by the total charge passed, i.e., nFE_cell, where:

• F = Faraday constant (96500 coulombs), corresponding to the charge of one mole of electrons.
• n = no. of moles of electrons flowing through the cell.
• E_cell = EMF or cell potential.

For a voltaic cell, the redox reaction is spontaneous. For a spontaneous process, the value of Gibbs energy change ΔG = −ve, and the decrease in Gibbs energy would be equal to the electrical work done by the cell, hence:

ΔG = −nFE_cell

Unit for F is Coulomb. Unit for E_cell is volt. Therefore, the unit for electric work is Coulomb × Volt = CV, and 1 Coulomb × 1 Volt = 1 Joule.

At standard state (1 M concentration of ionic species), 1 atm pressure, and 25 °C, the Gibbs energy change is known as the standard Gibbs energy change of a reaction (ΔG°), and the corresponding EMF will be the standard potential or EMF (E°_cell).

−ΔG° = nFE°_cell   or   ΔG° = −nFE°_cell

where ΔG = ΔH − TΔS (thermodynamic definition of Gibbs’ energy).

In everyday life, we use batteries to power our devices. A battery is typically an arrangement of one or more electrochemical cells connected in series. Commercial batteries must:

• Be lightweight and compact.
• Deliver a constant voltage over their useful lifetime.
• Be safe, reliable, and economical.
• Some batteries are rechargeable, and some are not.

Primary Cells	Secondary Cells
Non-rechargeable.	Rechargeable.
Cell reaction occurs in only one direction; it cannot be reversed.	Cell reaction is reversible; it can be recharged by passing electricity.
Become ‘dead’ after use.	Can be used repeatedly (hundreds of cycles).
Act as galvanic cells during discharge.	Act as galvanic cells during discharge and electrolytic cells during recharging.
Example: Dry cell (Leclanché cell)	Examples: Lead-acid battery, lithium-ion battery, nickel-cadmium battery.

1. Primary Cells (Non-rechargeable)

• In primary cells, the reaction occurs in only one direction and cannot be reversed.
• As a result, it cannot be charged or reused, but instead becomes dead after being used over a period of time.

For example, Dry cell, Fuel cell, Cadmium cell, etc.

Dry Cell (Leclanché cell / Zn–C battery)

The most common type of primary cell, a compact form of cell, was developed by a French engineer, George Leclanché (1866). Widely used in portable electric and electronic equipment. Each cell delivers a voltage of 1.5V.

Construction:

• Outer container: made up of Zn, which serves as an anode and provides electrons to the external circuit. The zinc is lined from the inside with an insulating paper.
• Central rod: A carbon rod, having a brass cap, acts as a cathode.
• Electrolyte paste: The space between cathode and anode is filled with a moist mixture of MnO₂, NH₄Cl (electrolyte), ZnCl₂ (as deliquescent), and charcoal powder.
• The porous paper lining prevents direct contact between the zinc container and the paste. This acts as a salt bridge.
• The cell is sealed from the top with wax.

The main reactions involved are:

At anode: Zn → Zn²⁺(aq) + 2e⁻

Zn²⁺ ions enter the electrolytic paste and react with NH₄Cl.

At cathode: MnO₂(s) + H₂O(l) + e⁻ → MnO(OH)(s) + OH⁻(aq)

2. Secondary Cells (Rechargeable)

• These can be recharged by passing electricity through them after every use; hence, they can be used again and again.
• In these cells, the electrical energy is stored in the form of chemical energy, hence called storage or accumulators.
• They act as galvanic cells during discharge and electrolytic cells during recharging.

For example: Lithium-ion battery, Acid lead storage battery, Nickel-Cadmium storage battery, etc.

The Acid Lead Storage Battery

A classic rechargeable (secondary) battery that has powered vehicles for over 150 years.

Construction:

• Anode: Grid of lead-antimony (Pb-Sb) alloy filled with spongy lead (Pb).
• Cathode: Grid of Pb-Sb alloy filled with lead(IV) oxide (PbO₂).
• Electrolyte: 20–38% sulphuric acid (H₂SO₄) solution.

• During discharge (when the battery powers a device): acts as a GALVANIC cell.
• During charging (when connected to a charger): acts as an ELECTROLYTIC cell.
• A fully charged battery can be safely stored for up to one year.
• Typical service life is about 3–10 years, depending on use and maintenance.

Conventional cells, such as dry cells and lead-acid batteries, have limitations including lower energy density, shorter lifespan, environmental concerns, and dependence on fossil-fuel-based electricity generation. Therefore, scientists have developed advanced electrochemical cells such as hydrogen–oxygen fuel cells and lithium-ion rechargeable batteries.

Lithium-ion Battery (Lithium-polymer)

The lithium-ion (Li-ion) battery is the most important rechargeable battery of our modern era. Since its commercial introduction in 1991, it has revolutionized portable electronics.

Why Lithium?

Lithium is the lightest metal (atomic mass = 6.94 g/mol) and has the most negative standard electrode potential (E° = −3.04 V), making it an excellent anode material with very high energy density.

Components:

Component	Material	Role
Anode (−)	Graphite intercalated with lithium (LiC₆)	Releases Li⁺ ions during discharge; stores Li during charging.
Cathode (+)	LiCoO₂, LiFePO₄, or LiNiMnCoO₂ (NMC)	Stores Li⁺ ions during discharge.
Electrolyte	Lithium salt (LiClO₄, LiBF₄) in organic solvent (e.g., ethylene carbonate)	Allows Li⁺ ion transport between electrodes.
Separator	Semi-permeable polymer membrane	Allows Li⁺ ions to pass but prevents electron flow (and short circuits).

In this cell, lithium undergoes oxidation to Li⁺ ions, which move from anode to cathode, where Li⁺ gets stored during discharge; the reaction reverses during charging. The semi-permeable barrier allows Li⁺ ions to move from anode to cathode and vice versa, but prohibits the flow of electrons inside the structure of the battery. Electrons are allowed to move across the external circuit, generating a cell potential of the magnitude 3.6V to 3.8V.

Cell Notation (typical):

Anode	Electrolyte	Cathode
LiC₆ (Lithiated)	Li–Polymer gel or LiClO₄ solution in ethylene carbonate	LiCoO₂ or LiFePO₄ or LiNiMnCoO₂

The reactions involved are:

At anode: LiC₆ − 1e⁻ → C₆ + Li⁺  (Oxidation)

At cathode: CoO₂ + Li⁺ + 1e⁻ → LiCoO₂  (Reduction)

Overall reaction: LiC₆ + CoO₂ ⇌ C₆ + LiCoO₂

• High energy density compared to other rechargeable batteries, allowing more energy storage in a smaller size.
• Lightweight and portable.
• Very low self-discharge rate: only 0.35–2.5% per month, retaining charge for longer periods.
• Charge–discharge efficiency: 80–90%.
• Cycle life: 400–1000 charge/discharge cycles. Can be recharged repeatedly, reducing waste and cost.
• Widely used in mobile phones, laptops, electric vehicles, and renewable energy storage systems.

Uses of Lithium-ion Battery:

1. In portable electronic devices such as laptops, digital cameras, mobile phones, and flashlights.
2. Medical equipment (pacemakers, hospital devices).
3. In electric vehicles (EVs).
4. Aerospace, military, and electric inverters, etc.

Fuel cells are different from batteries — they do not store energy internally. Instead, they convert fuel to electricity continuously, much like an engine, but with much higher efficiency and zero harmful emissions.

Hydrogen/Oxygen Fuel Cells

The most common and well-studied fuel cell is the hydrogen-oxygen fuel cell, which produces electricity by combining hydrogen fuel with oxygen, with water as the only product.

Construction and Working

• Hydrogen gas (H₂) and oxygen gas (O₂) are bubbled through porous carbon electrodes into a concentrated aqueous KOH or NaOH solution (electrolyte).
• Catalysts (usually platinum or nickel) are incorporated into the electrodes to speed up the reaction.
• H₂ is oxidised at the anode; O₂ is reduced at the cathode.

Reactions involved in this cell are:

At anode: [H₂(g) + 2 OH⁻(aq) → 2 H₂O(l) + 2 e⁻] × 2  (Oxidation half)

At cathode: O₂(g) + 2 H₂O(l) + 4 e⁻ → 4 OH⁻  (Reduction half)

Overall reaction: 2 H₂(g) + O₂(g) → 2 H₂O(l)

Energy conversion efficiency: ~60–70% (much better than a petrol engine at ~20–25%!)

Advantages of Hydrogen-Oxygen Fuel Cells

• Run continuously as long as fuel and oxygen are supplied.
• Significantly more efficient than conventional combustion engines.
• Pollution-free: the only product is pure water (H₂O).
• Silent operation — no moving parts.

Applications

• Transport: fuel cell buses and cars (e.g., Toyota Mirai).
• Stationary power generation for homes and industry.
• Space missions: NASA has used H₂–O₂ fuel cells since the Apollo programme.
• Emergency backup power systems.
• Medical devices like pacemakers.

Limitations

• Hydrogen is highly flammable and difficult to store safely in a vehicle.
• Large-scale, economical production of pure hydrogen remains a major challenge.
• Requires expensive platinum catalysts.