Electrochemistry: Complete Guide & Formula Sheet with Free PDF Download (JEE & NEET)
Electrochemistry — Competishun
Electrochemistry: Complete Guide & Formula Sheet with Free PDF Download (JEE & NEET)
Galvanic Cells · Nernst · Conductance · Faraday · Batteries
Electrochemistry is one of the most important and scoring chapters in physical chemistry. It carries high weightage in JEE and NEET, with 2-3 questions appearing every year.
Electrochemistry studies the interconversion of electrical and chemical energy. It covers galvanic cells, electrolytic cells, Nernst equation, conductance, Kohlrausch law, Faraday's laws of electrolysis, batteries, and corrosion.
This page gives you the complete Electrochemistry guide with all concepts explained in depth. You will find clear definitions, derivations, worked examples, and common mistakes to avoid. Download the free PDF below and keep it handy for quick revision before your JEE Main, JEE Advanced, or NEET exam.
E = E° - (0.0591/n) log QNernst Equation
Λ°m = ν+λ°+ + ν−λ°−Kohlrausch Law
w = ZItFaraday's Laws
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What is Electrochemistry?
Definition: Electrochemistry is the branch of chemistry that deals with the relationship between electrical energy and chemical changes. It studies the interconversion of electrical energy and chemical energy.
Electrochemistry is divided into two main branches:
Galvanic Cells
Convert chemical energy into electrical energy. Also called voltaic cells. Examples: Daniell cell, dry cell, lead storage battery.
Electrolytic Cells
Convert electrical energy into chemical energy. Uses external voltage to drive non-spontaneous reactions. Examples: electrolysis of water, electroplating.
Key Insight: Electrochemistry connects thermodynamics (ΔG), equilibrium (K), and electricity (E). The relationship ΔG° = -nFE° is the bridge between all three. Understanding this connection makes the entire chapter much easier to grasp.
Galvanic Cells — Electrochemistry
Definition: A galvanic cell (or voltaic cell) is a device that converts chemical energy into electrical energy through spontaneous redox reactions. The electrons flow from the anode to the cathode through an external circuit.
Components of a Galvanic Cell
- Anode (-): The electrode where oxidation occurs. Electrons are produced here and flow to the cathode.
- Cathode (+): The electrode where reduction occurs. Electrons are consumed here.
- Salt Bridge: Maintains electrical neutrality by allowing ions to flow. Usually contains KCl or KNO₃ in agar-agar gel.
- External Circuit: Allows electrons to flow from anode to cathode, producing an electric current.
The Daniell Cell — Classic Example
| Component | Details |
|---|---|
| Anode (-) | Zn electrode in ZnSO₄ solution. Zn → Zn²⁺ + 2e⁻ (Oxidation) |
| Cathode (+) | Cu electrode in CuSO₄ solution. Cu²⁺ + 2e⁻ → Cu (Reduction) |
| Salt Bridge | KCl or KNO₃ in agar-agar |
| Cell Notation | Zn(s) │ Zn²⁺(C₁) ‖ Cu²⁺(C₂) │ Cu(s) |
| Overall Reaction | Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) |
| E°cell | E°cell = E°cathode - E°anode = 0.34 - (-0.76) = 1.10 V |
| Electrons always flow from anode to cathode. Current flows from cathode to anode (opposite direction). | |
Important: In Electrochemistry, remember the mnemonic: "AN OX" (Anode = Oxidation) and "RED CAT" (Reduction = Cathode). This never changes, whether you are dealing with a galvanic cell or an electrolytic cell.
Standard Electrode Potential (E°)
The standard electrode potential is the potential of an electrode measured under standard conditions:
- Temperature: 298 K (25°C)
- Concentration: 1 M for all ions
- Pressure: 1 bar for gases
The Standard Hydrogen Electrode (SHE) is used as the reference electrode with E° = 0.000 V by convention.
Nernst Equation — Electrochemistry
Definition: The Nernst equation relates the electrode potential to the concentrations of ions in the cell reaction. It is one of the most important equations in Electrochemistry.
General Form
| Temperature | Formula |
|---|---|
| Any Temperature | E = E° - (RT/nF) ln Q |
| At 298 K (25°C) | E = E° - (0.0591/n) log Q |
| Where: E = cell potential, E° = standard cell potential, R = 8.314 J/mol·K, T = temperature, n = number of electrons transferred, F = 96500 C/mol, Q = reaction quotient. | |
Key Formula: E = E° - (0.0591/n) log Q. This is the most frequently used formula in Electrochemistry problems. Make sure you remember it for the exam.
Applications of Nernst Equation
- Calculation of cell potential under non-standard conditions — when concentrations are not 1 M
- Determination of equilibrium constant (K) — at equilibrium, E = 0, so log K = nE° / 0.0591
- Calculation of concentration of unknown ions — using concentration cells
- Calculation of pH of a solution — using hydrogen electrodes
Concentration Cells
A concentration cell has the same electrodes but different concentrations. For such cells, E°cell = 0.
| Type | Formula |
|---|---|
| Electrolyte Concentration Cell | E = (0.0591/n) log(C₂/C₁) where C₂ > C₁ |
| Gas Electrode Concentration Cell | E = (0.0591/n) log(P₁/P₂) |
| In concentration cells, the electrode in the more concentrated solution acts as the cathode. | |
Thermodynamics of Electrochemical Cells
In Electrochemistry, the cell potential is directly related to thermodynamic quantities.
| Quantity | Formula | Meaning |
|---|---|---|
| Gibbs Free Energy | ΔG° = -nFE° | Maximum work obtainable from the cell |
| Equilibrium Constant | log K = nE° / 0.0591 | Relates E° to K |
| Entropy Change | ΔS° = nF(∂E°/∂T) | Temperature dependence of E° |
| Enthalpy Change | ΔH° = ΔG° + TΔS° | From Gibbs-Helmholtz equation |
| If E°cell > 0, then ΔG° < 0 and the reaction is spontaneous. | ||
Key Insight: The relationship ΔG° = -nFE° is the bridge between Electrochemistry and Thermodynamics. This one equation connects three different topics in physical chemistry, so understanding it well can save you a lot of time in the exam.
Conductance in Electrolyte Solutions — Electrochemistry
Definition: Conductance is the ability of a solution to conduct electricity. It depends on the number of ions present and their mobility.
Important Terms and Definitions
| Quantity | Symbol | Formula | Unit |
|---|---|---|---|
| Resistance | R | R = ρ(l/A) | Ω (ohm) |
| Conductance | G | G = 1/R | S (siemens) or mho |
| Conductivity (Specific Conductance) | κ | κ = G·(l/A) = G·G* | S m⁻¹ or S cm⁻¹ |
| Cell Constant | G* | G* = l/A | m⁻¹ or cm⁻¹ |
| Molar Conductivity | Λm | Λm = κ/C = 1000κ/M | S m² mol⁻¹ |
| 1 S cm⁻¹ = 100 S m⁻¹. 1 S m² mol⁻¹ = 10⁴ S cm² mol⁻¹. These conversions are frequently tested. | |||
Effect of Dilution on Conductance
| Property | Strong Electrolyte | Weak Electrolyte |
|---|---|---|
| Conductivity (κ) | Decreases on dilution | Decreases on dilution |
| Molar Conductivity (Λm) | Increases slowly on dilution | Increases sharply on dilution |
| Reason | Ions move farther apart (mobility increases) | Degree of dissociation (α) increases |
| For strong electrolytes, Λm can be extrapolated to find Λ°m. For weak electrolytes, this is not possible. | ||
Important: The molar conductivity (Λm) of a solution is always higher when the solution is more dilute. This is because the ions have less interionic attraction and can move more freely.
Kohlrausch's Law — Electrochemistry
Definition: Kohlrausch's law states that the limiting molar conductivity (Λ°m) of an electrolyte is the sum of the individual contributions of its constituent ions, each multiplied by its stoichiometric coefficient.
Formula: Λ°m = ν+λ°+ + ν−λ°−
Where ν+ and ν− are the number of cations and anions per formula unit, and λ°+ and λ°− are their limiting ionic conductivities.
Where ν+ and ν− are the number of cations and anions per formula unit, and λ°+ and λ°− are their limiting ionic conductivities.
Applications of Kohlrausch's Law in Electrochemistry
- To find Λ°m of weak electrolytes: Example: Λ°(CH₃COOH) = Λ°(CH₃COONa) + Λ°(HCl) - Λ°(NaCl)
- To calculate the degree of dissociation (α): α = Λm / Λ°m
- To find the dissociation constant (Ka): Ka = Cα² / (1-α)
- To calculate solubility of sparingly soluble salts: S = 1000κ / Λ°m
Key Fact: λ°(H⁺) = 349.8 S cm² mol⁻¹ and λ°(OH⁻) = 198.5 S cm² mol⁻¹. These are much larger than other ions due to the Grotthuss mechanism (proton hopping). This is a frequently asked concept in Electrochemistry.
| Ion | λ° (S cm² mol⁻¹) | Explanation |
|---|---|---|
| H⁺ | 349.8 | Very high due to Grotthuss mechanism |
| OH⁻ | 198.5 | High due to Grotthuss mechanism |
| K⁺ | 73.5 | Normal ion mobility |
| Na⁺ | 50.1 | Lower than K⁺ due to larger hydrated size |
| Cl⁻ | 76.3 | Normal ion mobility |
| The order of ionic mobility: H⁺ > OH⁻ > K⁺ > Cl⁻ > Na⁺ > Li⁺ | ||
Faraday's Laws of Electrolysis — Electrochemistry
Definition: Faraday's laws describe the quantitative relationship between the amount of electricity passed through an electrolyte and the amount of substance deposited or liberated at the electrodes.
First Law of Faraday
The amount of substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed.
Formula: w = Z·I·t = (M·I·t)/(n·F)
Where: w = mass deposited, Z = electrochemical equivalent, I = current (A), t = time (s), M = molar mass, n = number of electrons, F = 96500 C/mol
Where: w = mass deposited, Z = electrochemical equivalent, I = current (A), t = time (s), M = molar mass, n = number of electrons, F = 96500 C/mol
Second Law of Faraday
When the same quantity of electricity is passed through different electrolytes in series, the amounts of substances deposited are proportional to their equivalent weights.
Formula: w₁/w₂ = E₁/E₂
Where E = equivalent weight = M/n
Where E = equivalent weight = M/n
Important Relations
- Moles of electrons: mol e⁻ = I·t / F
- Z (electrochemical equivalent): Z = E / F = M/(n·F)
- Volume of gas at STP: V = (I·t/F)·(22400/n) mL
Key Insight: In Electrochemistry, 1 Faraday = 96500 C = charge on 1 mole of electrons. This is the bridge between electricity and chemistry.
| Quantity | Formula | Example |
|---|---|---|
| Mass deposited | w = (M·I·t)/(n·F) | Cu (n=2): w = (63.5·I·t)/(2·96500) |
| Volume of gas | V = (I·t/F)·(22400/n) mL | H₂ (n=2): V = (I·t/F)·11200 mL |
| Current required | I = (w·n·F)/(M·t) | To deposit 1g Cu in 1h: I ≈ 0.84 A |
| In series, equivalents are equal: w₁/E₁ = w₂/E₂. This is a common JEE trap, so be careful. | ||
Electrolysis — Electrochemistry
Definition: Electrolysis is the process of decomposing an electrolyte by passing an electric current through it. It is the reverse of a galvanic cell.
Important Rules for Electrolysis
- At the Cathode (-): The ion with the highest reduction potential is discharged first.
- At the Anode (+): The ion with the lowest reduction potential (highest oxidation potential) is discharged first.
- Active Electrodes: If the anode is active (Cu, Ag, Ni), it dissolves instead of the electrolyte ions.
Preferential Discharge Series
| Electrode | Order of Discharge (Decreasing Priority) |
|---|---|
| Cathode (-) | Ag⁺ > Cu²⁺ > H⁺ (pH 0) > H⁺ (pH 7) > Zn²⁺ > H₂O > Na⁺ |
| Anode (+) | Cl⁻ > OH⁻ > SO₄²⁻ (or NO₃⁻, F⁻) |
| Never deposited from aqueous solution: Li⁺, K⁺, Ca²⁺, Na⁺, Mg²⁺, Al³⁺ → H₂ is produced instead. | |
Products of Electrolysis — Examples
| Electrolyte | Electrode | Cathode Product | Anode Product |
|---|---|---|---|
| Molten NaCl | Inert | Na metal | Cl₂ gas |
| Dilute NaCl (aq) | Pt | H₂ gas | O₂ gas |
| Conc. Brine (NaCl) | Pt | H₂ gas | Cl₂ gas |
| CuSO₄ (aq) | Pt | Cu metal | O₂ gas |
| CuSO₄ (aq) | Cu | Cu metal | Cu dissolves (active anode) |
| Active anode (Cu, Ag, Ni) dissolves itself, so the electrolyte concentration remains unchanged. | |||
Exception: O₂ has about 0.5 V overvoltage, so in concentrated brine, Cl⁻ (1.36 V) beats water (1.23 V), and Cl₂ is produced instead of O₂.
Batteries — Electrochemistry
Definition: A battery is a device that converts chemical energy into electrical energy. It consists of one or more electrochemical cells connected in series or parallel.
1. Lead Storage Battery (12 V)
| Component | Details |
|---|---|
| Anode (-) | Pb (lead) electrode |
| Cathode (+) | PbO₂ (lead dioxide) electrode |
| Electrolyte | 38% H₂SO₄ (density approximately 1.28 g/mL) |
| Discharge Reaction | Pb + PbO₂ + 2H₂SO₄ → 2PbSO₄ + 2H₂O |
| Charging Reaction | 2PbSO₄ + 2H₂O → Pb + PbO₂ + 2H₂SO₄ |
| On discharge, H₂SO₄ is used up, so density decreases. On charging, H₂SO₄ is regenerated, so density increases. | |
2. Dry Cell (1.5 V)
- Anode: Zn can (zinc container) → Zn → Zn²⁺ + 2e⁻
- Cathode: MnO₂ + C (graphite rod) → MnO₂ + NH₄⁺ + e⁻ → MnO(OH) + NH₃
- Electrolyte: NH₄Cl + ZnCl₂ paste
- Voltage: Approximately 1.5 V
3. Mercury Cell (1.35 V)
- Anode: Zn(Hg) (zinc amalgam) → Zn(Hg) + HgO → ZnO + Hg
- Cathode: HgO + C
- Electrolyte: KOH + ZnO
- Voltage: Constant 1.35 V (no ions in net reaction)
4. H₂-O₂ Fuel Cell
| Component | Details |
|---|---|
| Anode (-) | H₂ gas → 2H₂ + 4OH⁻ → 4H₂O + 4e⁻ |
| Cathode (+) | O₂ gas → O₂ + 2H₂O + 4e⁻ → 4OH⁻ |
| Overall Reaction | 2H₂ + O₂ → 2H₂O |
| E°cell | Approximately 1.23 V |
| Efficiency | η = (ΔG/ΔH) × 100 ≈ 70-75% |
| Fuel cells are efficient because they bypass thermal energy conversion. The only product is water. | |
Corrosion — Electrochemistry
Definition: Corrosion is the gradual destruction of metals by chemical or electrochemical reactions with their environment. Rusting of iron is the most common example.
Rusting of Iron — Electrochemical Explanation
| Electrode | Reaction |
|---|---|
| Anode | Fe → Fe²⁺ + 2e⁻ (E° = -0.44 V) |
| Cathode | O₂ + 4H⁺ + 4e⁻ → 2H₂O (E° = +1.23 V) |
| Overall E°cell | E°cell = 1.23 - (-0.44) = +1.67 V (spontaneous) |
| Rusting requires both O₂ and H₂O. Rust is Fe₂O₃·xH₂O, which is porous and flaky. | |
Key Difference: Rust (Fe₂O₃·xH₂O) is porous and flaky, so fresh metal keeps getting exposed. Al₂O₃ (aluminium oxide) is protective because it forms a thin, dense layer that stops further corrosion.
Methods to Prevent Corrosion
- Coating: Painting, oiling, greasing
- Galvanisation: Coating iron with zinc
- Cathodic Protection: Using a more active metal (sacrificial anode) like Mg or Zn
- Alloying: Making alloys like stainless steel (Fe + Cr + Ni)
All Electrochemistry Formulas at a Glance
| Formula | What It Means |
|---|---|
| E = E° - (0.0591/n) log Q | Nernst equation at 298 K |
| E°cell = E°cathode - E°anode | Standard cell potential |
| ΔG° = -nFE° | Gibbs free energy from cell potential |
| log K = nE°cell / 0.0591 | Equilibrium constant from cell potential |
| Λ°m = ν+λ°+ + ν−λ°− | Kohlrausch's law |
| α = Λm / Λ°m | Degree of dissociation |
| w = (M·I·t)/(n·F) | Faraday's first law |
| mol e⁻ = I·t / F | Moles of electrons |
| κ = G·G* | Conductivity from conductance |
| Λm = 1000κ/M | Molar conductivity |
| Memorise these formulas for Electrochemistry. They are the key to scoring full marks in this chapter. | |
Common Mistakes in Electrochemistry
- Forgetting to balance electrons in Nernst equation: Use n = number of electrons transferred in the balanced reaction.
- Using wrong sign for E°cell: E°cell = E°cathode - E°anode. Always use reduction potentials from the electrochemical series.
- Confusing anode and cathode in electrolytic cells: In electrolytic cells, anode is positive (+) and cathode is negative (-), which is opposite to galvanic cells.
- Forgetting that conductivity and molar conductivity change differently on dilution: κ decreases, Λm increases.
- Misapplying Faraday's laws: In series, equivalents are equal, not moles.
- Using incorrect units: 1 S cm⁻¹ = 100 S m⁻¹. This is a frequent source of errors.
Golden Rule: Always write the balanced half-reactions first. Identify the number of electrons (n) and use the correct sign conventions in Electrochemistry. This single habit will save you from most of the common mistakes.
Why Electrochemistry Matters for JEE and NEET
- High weightage: Electrochemistry appears in 2-3 questions in every JEE Main, JEE Advanced, and NEET chemistry paper.
- Foundation for physical chemistry: Understanding Electrochemistry helps you with thermodynamics, equilibrium, and chemical kinetics.
- Direct scoring: Many questions are direct formula-based, especially Nernst equation, conductance, and Faraday's laws.
- Real-world applications: Electrochemistry connects to batteries, corrosion, electroplating, and industrial electrolysis.
- Conceptual clarity: This chapter rewards students who understand the concepts rather than just memorizing formulas.
Why this guide helps: A comprehensive Electrochemistry guide with all concepts, definitions, derivations, and formulas saves you time during revision and helps you quickly recall everything during the exam. You won't need to look anywhere else.
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Frequently Asked Questions — Electrochemistry
What is Electrochemistry?
Electrochemistry is the study of the relationship between electrical energy and chemical changes. It covers galvanic cells that convert chemical energy to electrical energy, electrolytic cells that convert electrical energy to chemical energy, and the fundamental principles governing these processes including Nernst equation, conductance, and Faraday's laws.
What is the Nernst equation in Electrochemistry?
The Nernst equation in Electrochemistry is E = E° - (RT/nF) ln Q. At 298 K, it becomes E = E° - (0.0591/n) log Q. It relates the electrode potential to the concentrations of ions in the cell reaction and is used to calculate cell potential under non-standard conditions.
What is Kohlrausch's law of independent migration of ions?
Kohlrausch's law in Electrochemistry states that the limiting molar conductivity of an electrolyte is the sum of the contributions of its individual ions: Λ°m = ν+λ°+ + ν−λ°−. It is used to calculate the limiting molar conductivity of weak electrolytes, determine the degree of dissociation, and find the solubility of sparingly soluble salts.
What are Faraday's laws of electrolysis?
Faraday's laws in Electrochemistry state: First law: the amount of substance deposited or liberated at an electrode is directly proportional to the quantity of electricity passed (w = ZIt). Second law: when the same quantity of electricity is passed through different electrolytes, the amounts of substances deposited are proportional to their equivalent weights. 1 Faraday = 96500 Coulombs.
Can I download the Electrochemistry formula sheet PDF for free?
Yes. You can download the complete Electrochemistry formula sheet PDF for free using the download button on this page. It covers galvanic cells, Nernst equation, conductance, Kohlrausch law, Faraday's laws, batteries, and corrosion in one comprehensive place for quick revision before JEE and NEET exams.
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