Law of definite proportions

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In chemistry, the law of definite proportions, sometimes called Proust's Law, states that a chemical compound always contains exactly the same proportion of elements by mass. An equivalent statement is the law of constant composition, which states that all samples of a given chemical compound have the same elemental composition. For example, oxygen makes up 8/9 of the mass of any sample of pure water, while hydrogen makes up the remaining 1/9 of the mass. Along with the law of multiple proportions, the law of definite proportions forms the basis of stoichiometry.[1]

Contents

History

This observation was first made by the French chemist Joseph Proust, based on several experiments conducted between 1798 and 1804.[2] Based on such observations, Proust made statements like this one, in 1806:

The law of definite proportions might seem obvious to the modern chemist, inherent in the very definition of a chemical compound. At the end of the 18th century, however, when the concept of a chemical compound had not yet been fully developed, the law was novel. In fact, when first proposed, it was a controversial statement and was opposed by other chemists, most notably Proust's fellow Frenchman Claude Louis Berthollet, who argued that the elements could combine in any proportion.[3] The very existence of this debate underscores that at the time, the distinction between pure chemical compounds and mixtures had not yet been fully developed.[4]

The law of definite proportions contributed to, and was placed on a firm theoretical basis by, the atomic theory that John Dalton promoted beginning in 1803, which explained matter as consisting of discrete atoms, that there was one type of atom for each element, and that the compounds were made of combinations of different types of atoms in fixed proportions.[5]

A related early idea was Prout's hypothesis, which supposed that hydrogen was the only functional unit, and was related to the whole number rule, which was the rule of thumb that atomic masses were whole number multiples of the mass of hydrogen. This was later rejected in the 1820s and 30s following more refined measurements of atomic mass, notably by Jöns Jacob Berzelius, which revealed in particular that the atomic mass of chlorine was 35.45, which was incompatible with the hypothesis. Today however this discrepancy is understood (since the 1920s) to be due to the presence of isotopes, and the mass of given isotopes is very close to satisfying the whole number rule, with the mass defect caused by differing binding energies being significantly smaller.

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