Electrochemistry Formula Sheet PDF for JEE (N Awasthi Companion)

Here are three free PDFs for JEE electrochemistry: a formula sheet with the Nernst equation, Faraday’s laws, and conductance; a one-page cheat sheet with exam shortcuts; and a list of eleven common traps. These are our original sheets from the 1% Syndicate, built for fast, effective revision. All our PDFs are completely free to download.

What these PDFs are (and what they are not)

Let’s be clear. These aren’t scanned pages from a textbook or some random file scraped from the web. These are our own original sheets, created by the AllAboutIIT 1% Syndicate. We wrote them in our own words to be a strategic companion, not a replacement for your core study material. They’re designed to sit right next to a problem book like N Awasthi (N Avasthi) for physical chemistry, giving you the theoretical tools you need without the fluff.

These sheets are not the book itself. They don’t contain its text, problems, or solutions. They are our synthesis of what actually matters for solving problems and acing exams. They are free for your personal study. Don’t try to sell them.

Not affiliated with the author or publisher of any book mentioned.

Download the Electrochemistry PDFs

We’ve broken this down into three files. One for deep study, one for rapid revision, and one for damage control. Use them all.

Electrochemistry formula sheet (PDF)

This is the master document. It contains every important formula, definition, and relationship you need for the entire chapter, laid out logically. It’s your primary reference when you’re first learning the chapter or doing a deep revision session. Don’t just read it; understand the conditions under which each formula applies.

Faraday’s first and second laws of electrolysis
Standard cell potential (E°cell) calculation
The Nernst equation in its general form and at 298 K
The relationship between Gibbs free energy, cell potential, and the equilibrium constant (ΔG = -nFE)
The temperature coefficient of cell potential
Formulas for concentration cells
Conductance, resistivity, conductivity, cell constant, and molar conductivity
Kohlrausch’s law and its application for weak electrolytes

Download the Electrochemistry formula sheet (PDF)

Electrochemistry cheat sheet (PDF)

This is your one-page weapon for the days leading up to a test. We’ve compressed the entire chapter into the most high-yield points, shortcuts, and mental triggers. This isn’t for learning the first time. It’s for activating what you already know, fast. Keep it handy when you’re solving a mixed-concept mock test.

A note from us: Focus on the Nernst equation and conductance formulas.
Key idea summaries: galvanic vs. electrolytic cells, anode/cathode rules, spontaneity checks.
Effect of dilution on conductivity and molar conductivity.
Exam shortcuts, like using the 0.0591/n form of the Nernst term.
The correct way to combine half-cells (hint: it’s via ΔG).
A high-value “if you see this, think this” list to connect problem triggers to the right concepts.
Key differences in electrolysis of aqueous solutions vs. molten salts.
Quick formula for log K from E°cell.

Download the Electrochemistry cheat sheet (PDF)

Electrochemistry mistakes and traps (PDF)

Marks aren’t just gained; they’re also saved. This document lists the most common, mark-killing mistakes students make in electrochemistry. We explain why the mistake happens (the faulty logic) and how to build a system to avoid it forever. Review this after every test and after every problem-solving session where you got something wrong.

The classic trap of adding E° values directly for half-reactions.
Mixing up ln and log10, or the RT/nF and 0.0591/n terms.
Setting up the reaction quotient (Qc) incorrectly, especially by including solids or liquids.
Confusing anode/cathode signs between galvanic and electrolytic cells.
Forgetting the value of ‘n’ (moles of electrons) in the Nernst equation.
Confusing conductivity (κ) and molar conductivity (Λm).
Unit conversion errors in conductance calculations.
Ignoring overpotential in the electrolysis of aqueous solutions.

Download the Electrochemistry mistakes and traps (PDF)

The chapter ideas in plain words

Forget the textbook jargon. Here’s what the core ideas really mean.

Galvanic vs. Electrolytic Cells

A galvanic (or voltaic) cell is a battery. It uses a spontaneous chemical reaction (ΔG < 0) to generate electricity (Ecell > 0). An electrolytic cell does the opposite. It uses electricity to force a non-spontaneous reaction to happen (ΔG > 0, Ecell < 0). Remember: In both cells, oxidation happens at the anode and reduction at the cathode (An Ox, Red Cat). The sign convention is the only thing that flips: in a galvanic cell, the anode is negative; in an electrolytic cell, it’s positive.

The Nernst Equation

This is the most important equation in the chapter. It connects the standard cell potential (E°cell), which is a fixed value for a given reaction at 25°C and 1M concentrations, to the actual cell potential (Ecell) under any other conditions of concentration and temperature. It tells you how the voltage changes as the battery runs down (i.e., as reactant concentrations decrease and product concentrations increase).

Ecell = E°cell - (RT/nF)lnQc

This holds under any condition. R is the gas constant (8.314 J/mol·K), T is temperature in Kelvin, n is moles of electrons transferred, F is Faraday’s constant (~96500 C/mol), and Qc is the reaction quotient.

Gibbs Free Energy and EMF

This is the bridge between thermodynamics and electrochemistry. The relationship ΔG = -nFEcell tells you if a reaction is spontaneous. Since n and F are positive, a positive cell potential (Ecell > 0) means a negative Gibbs free energy (ΔG < 0), which is the definition of a spontaneous process. This is why galvanic cells work. The standard version links the standard potential to the equilibrium constant, K.

ΔG° = -nFE°cell = -RTlnKc

This holds for standard conditions and at equilibrium.

Faraday’s Laws of Electrolysis

This is the quantitative part of electrolysis. The first law states that the mass of a substance deposited or liberated at an electrode is directly proportional to the quantity of charge passed through the electrolyte. The second law says that when the same quantity of charge is passed through different electrolytes, the masses of substances liberated are proportional to their chemical equivalent weights.

w = (M/nF) × I × t

Here, w is mass deposited, M is molar mass, n is the number of electrons in the half-reaction (n-factor), I is current in Amperes, and t is time in seconds.

Conductance and Conductivity

These terms describe how well a solution conducts electricity. Resistance (R) is the opposition to current flow. Conductance (G) is the inverse of resistance (G = 1/R). Conductivity, or specific conductance (κ, kappa), is the conductance of a 1 m³ (or 1 cm³) cube of the solution. It’s related to resistance by the cell’s geometry (l/A, the cell constant).

κ = (1/R) × (l/A)

Conductivity (κ) decreases on dilution because the number of ions per unit volume decreases.

Molar Conductivity

Molar conductivity (Λm, lambda-m) is a more useful measure because it accounts for concentration. It’s the conductivity produced by all the ions from one mole of electrolyte. It’s related to conductivity by Λm = κ / C (if C is in mol/m³) or Λm = (1000 × κ) / C (if C is in mol/L). Molar conductivity (Λm) always increases on dilution for both strong and weak electrolytes because ions get farther apart and inter-ionic attraction decreases.

Kohlrausch’s Law

This law is a lifesaver for weak electrolytes. It states that at infinite dilution, the molar conductivity of an electrolyte is the sum of the individual contributions of its cations and anions. You can’t measure Λ°m for a weak electrolyte directly by extrapolation, but you can calculate it by combining the Λ°m values of strong electrolytes.

Λ°m(CH3COOH) = Λ°m(CH3COONa) + Λ°m(HCl) - Λ°m(NaCl)

This works because you are effectively just adding the ionic conductivities of CH3COO- and H+.

Which formula, when

Use this table to map a problem’s situation to the right tool.

Situation	Formula to use	Watch out for
Finding cell potential at 298 K with non-standard concentrations	`Ecell = E°cell - (0.0591/n)log10(Qc)`	Make sure Qc = [Products]/[Reactants] is set up correctly. Don’t include solids.
Calculating mass deposited during electrolysis	`w = (M/nF) × I × t`	Units: I in Amperes, t in seconds. `n` is the n-factor for the half-reaction.
Checking if a cell reaction is spontaneous	`ΔG = -nFEcell`	Spontaneous means ΔG < 0, which requires Ecell > 0.
Finding the equilibrium constant from standard potential	`E°cell = (0.0591/n)log10(Kc)` (at 298 K)	This formula is only valid at equilibrium, where Ecell = 0.
Combining two half-reactions to find a third E°	`ΔG°3 = ΔG°1 + ΔG°2`, then `E°3 = -ΔG°3 / (n3 × F)`	Never add E° values directly. E° is intensive, ΔG° is extensive.
A concentration cell (same electrodes, different concentrations)	`Ecell = -(0.0591/n)log10(C_anode / C_cathode)`	The lower concentration side is always the anode.
Finding the degree of dissociation (α) of a weak electrolyte	`α = Λm / Λ°m`	Find Λm from experimental data (`κ/C`) and Λ°m using Kohlrausch’s law.
Relating resistance, conductivity, and cell constant (G*)	`κ = (1/R) × G` where `G = l/A`	The cell constant G* is specific to the physical cell and is often given.

Mistakes that cost the most marks

These aren’t just careless errors. They are fundamental misunderstandings that will cost you marks every single time.

Adding E° values directly when combining half-reactions.
This is the number one trap. E° is an intensive property (like density), you can’t add it. Gibbs Free Energy (ΔG°) is extensive (like mass). Convert each E° to ΔG° using ΔG° = -nFE°, add the ΔG° values, and then convert the final ΔG° back to the new E°.

Setting up the reaction quotient (Qc) incorrectly.
Students forget that Qc is [Products]^p / [Reactants]^r. They flip it, or worse, they include the concentrations of pure solids or liquids, which should be taken as 1. Always write the balanced cell reaction first before writing the Qc expression.

Confusing the anode and cathode signs.
Remember this system: In a Galvanic cell, the anode is Negative (it’s the source of electrons). In an Electrolytic cell, the anode is Positive (it’s where electrons are pulled away). Oxidation always happens at the anode, reduction at the cathode. No exceptions.

Forgetting the ‘n’ in the Nernst and Faraday equations.
‘n’ is the number of moles of electrons transferred in the balanced redox reaction. Forgetting it or getting it wrong (e.g., using n=1 for the Cu^2+ + 2e- -> Cu reaction) throws off the entire calculation. Always balance the half-reactions to find ‘n’.

Mixing up conductivity (κ) and molar conductivity (Λm).
Conductivity (κ) is about the solution in a 1 cm³ box; it decreases with dilution. Molar conductivity (Λm) is about 1 mole of electrolyte, no matter the volume; it increases with dilution. Know the difference and which one the question is asking for.

Forgetting stoichiometric coefficients in Kohlrausch’s law.
For an electrolyte like MgCl2, the law is Λ°m(MgCl2) = λ°(Mg^2+) + 2 × λ°(Cl-). Students often forget the ‘2’ for the chloride ion. Pay attention to the stoichiometry of the salt.

Two worked examples (our own)

No book problems. Just simple, clear examples to show the system.

Example 1: Nernst Equation Calculation

Given: A galvanic cell is constructed at 298 K with a Zinc electrode in 0.01 M ZnSO4 and a Copper electrode in 0.1 M CuSO4. Given E°(Zn^2+/Zn) = -0.76 V and E°(Cu^2+/Cu) = +0.34 V. Calculate the cell potential.

Method:

Identify Anode/Cathode: The electrode with the lower reduction potential is the anode. Zinc (-0.76 V) is lower than Copper (+0.34 V). So, Zn is the anode (oxidation) and Cu is the cathode (reduction).
Calculate E°cell: E°cell = E°cathode - E°anode = (+0.34 V) - (-0.76 V) = 1.10 V.
Write the cell reaction and find n:
Anode: Zn(s) -> Zn^2+(aq) + 2e-
Cathode: Cu^2+(aq) + 2e- -> Cu(s)
Overall: Zn(s) + Cu^2+(aq) -> Zn^2+(aq) + Cu(s).
The number of electrons transferred, n, is 2.
Set up the reaction quotient, Qc: Qc = [Products]/[Reactants] = [Zn^2+]/[Cu^2+] = 0.01 / 0.1 = 0.1.
Apply the Nernst Equation:
Ecell = E°cell - (0.0591/n)log10(Qc)
Ecell = 1.10 - (0.0591/2)log10(0.1)
Ecell = 1.10 - (0.02955) × (-1)
Ecell = 1.10 + 0.02955

Answer: Ecell = 1.12955 V

Example 2: Faraday’s Law Calculation

Given: A current of 5 Amperes is passed through a solution of molten Al2O3 for 20 minutes. Calculate the mass of Aluminum metal deposited at the cathode. (Molar mass of Al = 27 g/mol, 1 F = 96500 C/mol).

Method:

Identify the half-reaction and find n: Aluminum is deposited from Al^3+ ions.
Cathode reaction: Al^3+ + 3e- -> Al(s).
The number of electrons, n, is 3.
Calculate the total charge (Q) passed:
Current I = 5 A.
Time t = 20 minutes = 20 × 60 = 1200 seconds.
Q = I × t = 5 A × 1200 s = 6000 Coulombs.
Apply Faraday’s Law formula: The mass deposited w is given by w = (M/nF) × Q.
w = (27 / (3 × 96500)) × 6000
w = (9 / 96500) × 6000
w = 54000 / 96500

Answer: w ≈ 0.56 g

How to use these sheets with a problem book

Don’t just hoard the PDFs. Integrate them into your workflow. Here’s a simple system that works.

Prime Your Brain: Before you start solving problems from a book like N Awasthi, spend 15 minutes reading our Formula Sheet. Don’t memorize. Just get familiar with the tools and the structure of the chapter.
Attempt Problems Blind: Dive into the exercises. Try to solve a block of problems without looking at the solutions or our sheets. Struggle is good; it’s where learning happens.
Analyze Your Mistakes: For every problem you get wrong or couldn’t solve, stop. First, check our Mistakes and Traps sheet. Is your error listed there? Understand the faulty logic. Then, use the Formula Sheet to find the correct principle and re-solve the problem.
Consolidate and Revise: Before a test or moving to the next chapter, use the Cheat Sheet for a rapid 20-minute revision. This will solidify the main ideas and shortcuts in your mind.

A three-session plan for electrochemistry

This isn’t about studying 16 hours a day. It’s about focused, 60-90 minute sessions.

Session 1: Galvanic Cells & Thermodynamics. Your only goal is to master the galvanic cell. Focus on identifying anode/cathode, calculating E°cell, applying the Nernst equation, and understanding the link to ΔG and spontaneity. Solve problems only on this topic.
Session 2: Conductance. Forget cells for now. Focus entirely on the second half of the chapter. Understand the definitions of R, G, κ, Λm. Master the effect of dilution. Solve problems involving cell constant, calculation of Λm, and Kohlrausch’s law for weak electrolytes.
Session 3: Electrolysis & Consolidation. Focus on the quantitative aspects. Master Faraday’s laws. Solve problems calculating mass, volume of gas, or time/current required. End the session by solving a few mixed problems that combine concepts from all three sessions.

Quick self-test

Test your core understanding. No calculators needed for these.

A working galvanic cell has Ecell = +0.80 V. What is the sign of ΔG for the cell reaction?
In a concentration cell with Zn electrodes and ZnSO4 solutions of 0.01 M and 1.0 M, which side is the anode?
A current is passed through concentrated aqueous NaCl (brine) using inert electrodes. Which gas is produced at the anode in practice, and why not oxygen?
If the molar conductivity of a weak acid (Λm) is 15 and its molar conductivity at infinite dilution (Λ°m) is 300 (in the same units), what is its degree of dissociation (α)?
How many Faradays of charge are required to reduce 1 mole of MnO4- to Mn^2+ in an acidic medium?

Answers: 1. Negative (since ΔG = -nFEcell). 2. The 0.01 M side (lower concentration). 3. Cl2 gas. Thermodynamically water is slightly easier to oxidise, but oxygen needs a large overpotential, so Cl- is oxidised instead (the chlor-alkali process). 4. α = Λm / Λ°m = 15/300 = 0.05 or 5%. 5. 5 Faradays (The oxidation state of Mn changes from +7 to +2, a change of 5).

FAQ

What is the Nernst equation at 298 K?
It’s a simplified version for standard room temperature (25°C or 298 K). By plugging in the values for R, T, and F, and converting the natural log (ln) to base-10 log (log10), the term (RT/F)ln(Qc) becomes (0.0591)log10(Qc). The equation is Ecell = E°cell - (0.0591/n)log10(Qc).

Why do we add ΔG and not E when combining half-cell potentials?
Electrode potential (E) is an intensive property, meaning it doesn’t depend on the amount of substance. Gibbs free energy (ΔG) is an extensive property, meaning it does. You can’t add intensive properties (the density of two gold blocks isn’t their combined density). You can add extensive properties (the mass of two gold blocks is their combined mass). So, you convert E to the extensive property ΔG, add the ΔG values, then convert back.

What happens to conductivity and molar conductivity on dilution?
On dilution (adding more solvent), conductivity (κ) decreases. This is because the number of ions per unit volume decreases. However, molar conductivity (Λm) increases. This is because the total volume for one mole of electrolyte increases, giving ions more freedom to move and reducing inter-ionic forces.

How does Kohlrausch’s law help for weak electrolytes?
You can’t determine the molar conductivity at infinite dilution (Λ°m) for a weak electrolyte by just extrapolating a graph, because it never fully dissociates. Kohlrausch’s law lets you calculate it by cleverly adding and subtracting the known Λ°m values of strong electrolytes that contain the ions of your weak electrolyte.

Are these PDFs free to download?
Yes. All our resource sheets, including the formula sheet, cheat sheet, and mistakes list for every chapter, are completely free for personal use. The mission of the 1% Syndicate is to provide strategic tools, not to nickel-and-dime students.

What is the difference between Ecell and E°cell?
E°cell is the Standard Cell Potential. This is the potential of a cell measured under standard conditions: all solutions at 1 M concentration, all gases at 1 atm pressure, and the temperature at 298 K (25°C). Ecell is the Cell Potential under any non-standard set of conditions. The Nernst equation is what connects the two.

What is the role of the salt bridge?
The salt bridge has two critical jobs. First, it completes the electrical circuit by allowing ions to flow between the two half-cells. Second, it maintains electrical neutrality in each half-cell by migrating its ions to counteract the charge buildup from the electrode reactions. Without it, the cell would stop working almost instantly.

Once you’ve mastered the system for Electrochemistry, apply the same strategic thinking to other critical chapters.

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