For a binary solution in equilibrium with its own vapor, the Phase Rule tells us that the system possesses two degrees of freedom. So if temperature is held constant, there will be a relation at equilibrium between pressure and composition, corresponding to the observed reduction of vapor pressure by solutes. And if the solution is in equilibrium with solid solvent, there will be a relation at constant pressure between temperature and composition, corresponding to the observed depression of freezing point by solutes.
Back in the 1870s, before the thermodynamics of colligative properties had been placed on a theoretical footing, these relations had been discovered in Grenoble, France, by physicist François-Marie Raoult in connexion with his work on solutions, which occupied the last two decades of his life.
Raoult’s first paper describing solute-mediated depression of freezing point was published in 1878. It was a fertile area of study which Raoult appears to have had to himself, and not surprisingly he was the first to discover empirical relationships between quantities. One such relation he found was between the depression of freezing point and the molality of the solute:
Freezing-point depression = ΔTfp = Kfp x msolute
where Kfp is the cryoscopic constant for the solvent in degrees per molal. Raoult conducted many measurements of solutes in various solvents, and over time built up a body of data in support of the above relation. But in 1884, Raoult discovered a curious exception to the rule when he used sodium chloride as a solute – the effect on the freezing point of water was nearly twice as large as it should be. There was something peculiar about the behavior of common salt in solution.
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In the same year that Raoult in France discovered the anomalous colligative actions of sodium chloride as a solute, a 24-year-old student in Sweden named Svante Arrhenius submitted to Uppsala University a 150-page doctoral dissertation on electrolytic conductivity in aqueous solution. In it, he advanced the thesis that the conductivity of solutions of salts in water was due to the dissociation of the salt into oppositely charged particles, to which Faraday had given the name “ions” fifty years earlier. But while Faraday believed that ions were only produced during electrolysis, Arrhenius asserted that even in the absence of an electric current, solutions of salts contained ions. And he went as far as proposing that chemical reactions in solution were reactions between ions.
The professors at Uppsala were incredulous, and duly gave Arrhenius and his far-fetched ideas the minimum mark. But they underestimated the self-belief and resolve of their young student. Arrhenius followed up on their rebuttal by sending his dissertation to cutting-edge figures in Europe such as Wilhelm Ostwald and Jacobus Henricus van ‘t Hoff, who were actively developing the new science of physical chemistry. They were far more impressed, and following the award of a travel grant from the Swedish Academy of Sciences, Arrhenius was able to study with Ostwald in Riga, Kohlrausch in Germany, Boltzmann in Austria, and van ‘t Hoff in Amsterdam.
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Raoult’s discovery of the anomalous effect of common salt on the freezing point of water attracted the interest of JH van ‘t Hoff, who in 1887 subjected this peculiar result to detailed study. Investigating a number of ‘misbehaving’ salts, van ‘t Hoff found in each instance that the ratio of the measured freezing-point depression to the expected value approached a whole number as the solutions became increasingly dilute.
In the case of common salt, sodium chloride, the limiting ratio was 2. For sodium sulfate on the other hand, the ratio was 3, and for aluminum sulfate it was 5.
At the time when van ‘t Hoff found these whole number relations, Svante Arrhenius just happened to be visiting Amsterdam on his study tour. Arrhenius saw in these results the affirmation of his doctoral thesis, and could immediately supply the explanation. In aqueous solution, sodium chloride dissociates into sodium and chloride ions, so there are really two sets of solutes. Thus the total molality of the fully dissociated (ionised) solute will be double its undissociated value, and the freezing-point depression will be twice the expected amount.
By parallel reasoning, sodium sulfate dissociates into 3 aqueous ions (2 sodium ions and one sulfate ion) and in the case of aluminum sulfate there are 5 ions (2 aluminum ions and 3 sulfate ions).
It was compelling logic; the truth of Arrhenius’ thesis of ions in solution, and its implications for the understanding of chemical reactions and bonding, could no longer be denied.
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François-Marie Raoult’s work on solutions, and his discovery of the uncommon effect of common salt on the depression of freezing point, marked the start of a chain of circumstances that directly contributed to the founding of physical chemistry as a modern science. Not only did it provide affirmation of electrolytic dissociation and the existence of ions in solution, it also brought together the bright minds of Svante Arrhenius and Jacobus Henricus van ‘t Hoff, who with Wilhelm Ostwald were to propel physical chemistry into the modern age.
The Nobel Prize in Chemistry 1901 was awarded to JH van ‘t Hoff “in recognition of the extraordinary services he has rendered by the discovery of the laws of chemical dynamics and osmotic pressure in solutions”.
The Nobel Prize in Chemistry 1903 was awarded to Svante Arrhenius “in recognition of the extraordinary services he has rendered to the advancement of chemistry by his electrolytic theory of dissociation”.
The Nobel Prize in Chemistry 1909 was awarded to Wilhelm Ostwald “in recognition of his work on catalysis and for his investigations into the fundamental principles governing chemical equilibria and rates of reaction”.
photo credits: nobelprize.org