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Your chemistry book treats acetic acid and potassium chloride differently. Acetic acid is a weak electrolyte, only partly dissociated. Potassium chloride is a strong electrolyte, almost completely ionised. That second statement was not obvious a century ago, and one of the people who argued for it was a chemist in Calcutta. This post covers J. C. Ghosh's 1918 theory, what it got right, and how to handle the conductivity calculations in JEE and NEET chemistry.
Who was J. C. Ghosh?
The Government of India's notice on his death, in the Gazette of India of 23 January 1959, summarises his career. Dr. Jnan Chandra Ghosh died at Calcutta on 21 January 1959. He started in 1915 as Lecturer in Chemistry at Calcutta University, then served at Dacca University from 1921 to 1939 as Professor and Head of Chemistry. From 1939 to 1947 he was Director of the Indian Institute of Science, Bangalore. After Independence he became Director-General of Industries and Supplies, and later the first Director of IIT Kharagpur, Vice-Chancellor of Calcutta University in 1954 to 1955, and from 1955 a Member of the Planning Commission. The notice adds that he was President of the Indian Science Congress in 1939 and helped establish the Indian Chemical Society.
What problem was he trying to solve?
In 1918 Ghosh published a series of papers in the Journal of the Chemical Society, beginning with The abnormality of strong electrolytes. Part I, sent from the Physico-Chemical Laboratory, University College of Science, Calcutta. The full text is in the 1918 volume on the Internet Archive.
He set out the puzzle in the first paragraph. Arrhenius had proposed that a salt in water is in equilibrium between ions and undissociated molecules, with a degree of dissociation α. Ostwald applied the law of mass action to this equilibrium. Ghosh wrote that Ostwald's dilution law "fails absolutely in the case of strong electrolytes," and that none of the proposed fixes, such as hydration of ions or complex formation, had been satisfactory.
What did Ghosh propose?
His answer was to drop the equilibrium altogether. In his words: "In solutions of strong electrolytes only ions are present, and the attractive forces between ions are only governed by the physical laws of electrostatic attraction."
The rest of the paper builds a model on that idea:
Ions arrange themselves like a loose crystal. He assumed the average arrangement of ions in solution resembles the arrangement of atoms in the solid salt.
Only free ions carry current. An ion conducts only if its kinetic energy is enough to escape the pull of its neighbours. He used the kinetic theory of gases to estimate what fraction of ions can do this.
At infinite dilution, all ions are free. The ions are too far apart to attract each other, so molar conductivity reaches its maximum.
So, for Ghosh, the ratio of molar conductivity at a given concentration to its value at infinite dilution was not a degree of dissociation. It measured the fraction of ions free to move. He compared his calculated values with measured conductivities of salts and reported that, in one series, "The difference between the observed and the calculated values is rarely greater than 1 per cent."
His particular formula is not the one in your textbook today, but the starting point you now find in NCERT is the one he argued for. The NCERT Class 11 chapter Equilibrium says strong electrolytes "are ionized almost completely," and the Class 12 chapter Electrochemistry explains their conductivity changes without any dissociation equilibrium.
Worked example: Kohlrausch and the √c plot for KCl
NCERT gives two tools. For a strong electrolyte, Λm = Λ°m - A√c. And by Kohlrausch's law, Λ°m is the sum of the limiting molar conductivities of the ions. From NCERT's Table 2.4 at 298 K, λ°(K+) = 73.5 and λ°(Cl-) = 76.3 S cm² mol⁻¹.
Limiting molar conductivity from ions. Λ°m(KCl) = 73.5 + 76.3 = 149.8 S cm² mol⁻¹.
Now take two sample readings (made-up numbers close to real KCl data): Λm = 146.8 at c = 0.0010 mol L⁻¹, and Λm = 140.4 at c = 0.0100 mol L⁻¹.
Square roots. √0.0010 = 0.0316; √0.0100 = 0.100.
Slope. A = (146.8 - 140.4) / (0.100 - 0.0316) = 6.4 / 0.0684 = 93.6.
Intercept. Λ°m = 140.4 + 93.6 x 0.100 = 149.8 S cm² mol⁻¹, the same as step 1. The two methods agree.
The ratio Ghosh cared about. At 0.0100 mol L⁻¹, Λm / Λ°m = 140.4 / 149.8 = 0.937.
For acetic acid you would call that ratio α. For KCl, in Ghosh's picture, every ion is present, but only about 94 per cent are free to carry current at that concentration; the rest are held back by their neighbours.
Which mistakes cost marks?
Extrapolating a weak electrolyte. The √c straight line works for strong electrolytes. For acetic acid, Λm rises steeply at low concentration, so NCERT obtains Λ°m from Kohlrausch's law instead.
Mixing up units. S cm² mol⁻¹ and S m² mol⁻¹ differ by a factor of 10⁴. Convert before combining numbers from different tables.
Using c in the wrong place. Plot Λm against √c, not c, and keep c in mol L⁻¹.
Our post on conductance and Kohlrausch's law has more practice. The JEE (Main) 2026 syllabus lists weak and strong electrolytes and Kohlrausch's law; check the current year's syllabus and NEET-UG bulletin before you plan.
What can a student take from Ghosh's way of working?
Question the assumption, not just the formula. Others tried to patch Ostwald's law. Ghosh asked whether the equilibrium existed at all.
Make the theory give numbers. He compared calculated and measured conductivities and said how close they were.
Use science to build institutions. The Gazette credits him with laying sound foundations for higher technical education in India.
When a conductivity problem goes wrong at the set-up, the Fermi Problem Library has worked problems with a step-by-step solver that explains the reasoning rather than printing an answer.
Try this now: use NCERT's Table 2.4 to find Λ°m for NaCl, then explain in one sentence why the √c method would fail for acetic acid.
