The law of definite proportions, first explained in the late 1700s by the chemist Joseph Proust, is the foundation for modern science's understanding of chemical combinations. It says that, in any volume or mass, the elements of a chemical compound will maintain their set proportion. For example, a commonly known chemical compound is pure water, made up of hydrogen and oxygen in the formula H2O. The law of definite proportions says that regardless of the quantity of water — whether a glass, a rain barrel, or an eyedropper — the ratio of hydrogen to oxygen will always be one part hydrogen to eight parts oxygen. This law applies to the proportions of almost all chemical compounds.
Proust discovered the law while conducting experiments to determine the formulas of chemical compounds. His experiments over a six-year period were initially on metal compounds, and his conclusions differed from the established science of the day. Proust's discoveries were strongly contested by other scientists. It is believed that this reaction was due to confusion on the part of most 18th century scientists about the differences between pure and mixed chemical compounds.
One scientist who did not disagree with Proust was John Dalton, who at the same time was developing his theory of the law of multiple proportions. Coming at the principle from a different avenue, he noticed that when compounds were made using differing methods, their ratios were directly proportionate to the original compound elements. Further, he claimed these ratios were always expressed as whole numbers. When he heard Proust’s law of definite proportions, he realized that this law, combined with the law of multiple proportions, formed the basis of the earliest atomic theory, which explained the behavior of atoms according to fixed laws.
Today, scientists consider the law of definite proportions a critical scientific discovery. It is not, however, universally true. There are some chemical compounds that combine outside of the strict proportions of this law. In the 18th century, experimentation was not as precise as it would become in later centuries; measurements were not reported with enough exactitude to notice variations among the elements known at that time. Additionally, isotopes and their influences upon compounds had not yet been discovered. Factoring the impact of light and heavy isotopes into the analysis of atomic weights can account for the exceptions to the rule.
