In chemistry, the Law of Definite Proportion, also known as Proust's Law or the Law of Constant Composition, specifies that a chemical compound's constituent elements are usually in a specified mass ratio, regardless of its source or preparation process. We can also state the law of definite proportion in a different manner as well.
Example of Law of definite proportions
For example, oxygen makes up nearly 8/9 of any sample of pure water mass, while hydrogen makes up the rest of the 1/9 mass: the mass of two elements present in a compound are in the same ratio always. Along with the Law of multiple proportions, the Law of definite proportions produces the stoichiometry basis.
The Law of constant proportion was given in 1797 by Joseph Proust. This particular observation was made first by the Chemist and English theologian named Joseph Priestley and Antoine Lavoisier, who is a French nobleman and the chemist centered on the combustion process.
We may infer from these tests that, like many other metals, iron is subject to the Law of nature that governs each true combination, that is, it unites with two constant oxygen proportions, as mentioned at the beginning of this memoir. In this respect, it does not vary from mercury, lead, and tin, and, in a word, almost with every known combustible.
This Law of definite proportions can seem obvious to the modern chemist, which is inherent in the chemical compound definition. However, at the end of the 18 century, when the chemical compound concept had not yet been fully developed, the particular Law was novel. In fact, when this Law was first formulated, it was a divisive argument, with other chemists, most notably Proust's French colleague Claude Louis Berthollet, arguing that the elements could mix in some proportion. This debate demonstrates that the distinction between pure chemical compounds and their mixtures at the time had not yet been fully developed.
The Law of definite proportions was being contributed to and was placed on the basis of firm theoretic by the atomic theory that beginning in 1803, John Dalton promoted has explained the matter as holding of discrete atoms that there exists one type of atom for every element and that the compounds were prepared of combinations of various atom types in the fixed proportions.
One of the early related ideas was Prout's hypothesis, which was formulated by an English chemist William Prout, who has proposed that the hydrogen atom was the fundamental atomic unit. From this, the hypothesis was derived from the whole number rule - the rule of thumb that atomic masses were the whole number multiples of hydrogen mass.
Later, this was rejected in the 1820s and 1830s following atomic masses more refined measurements, especially by Jöns Jacob Berzelius, that revealed in specific that the atomic mass of chlorine has 35.45, which was incompatible with the hypothesis. The inclusion of isotopes has been used to describe this difference since the 1920s; every atomic isotope mass is closer to meeting the whole number law, with the mass fault created by differing binding energies being slightly lower.
Although it is very useful in modern chemistry foundation, universally, the Law of definite proportions is not true. There exist the non-stoichiometric compounds whose elemental composition can differ from one to another sample. Such types of compounds follow the Law of multiple proportions. For example, the iron oxide wüstite, which may hold between 0.83 and 0.95 iron atoms for each of the oxygen atoms, and therefore contain anywhere between a percentage of 23 and 25 oxygen by mass. The ideal formula is given as FeO, but because of the crystallographic vacancies, it is up to FeO.95O. In general, Proust's measurements were not accurate enough to spot such differences.
Also, the element's isotopic composition can differ based on its source; thus, its contribution to even a pure stoichiometric compound mass can differ. This variation can be used in radiometric dating since atmospheric, astronomical, crustal, oceanic, and deep Earth processes can concentrate few environmental isotopes preferentially. With the hydrogen and its isotope exception, usually, the effect is small, but it is measurable with modern-day instrumentation.
As an additional note, several natural polymers differ in composition (for example DNA, carbohydrates, and proteins) even when "pure". Generally, polymers are not considered as the "pure chemical compounds" except when their molecular weight is uniform (which is mono-disperse), and their stoichiometry is constant. In these unusual cases, they still can violate the Law because of the isotopic variations.