Science, Tech, Math › Science Law of Constant Composition in Chemistry Understanding the Mass Ratio Between Elements Share Flipboard Email Print According to the Law of Constant Composition, all samples of a compound contain the same mass ratios of atoms of the elements. Rafe Swan / Getty Images Science Chemistry Chemical Laws Basics Molecules Periodic Table Projects & Experiments Scientific Method Biochemistry Physical Chemistry Medical Chemistry Chemistry In Everyday Life Famous Chemists Activities for Kids Abbreviations & Acronyms Biology Physics Geology Astronomy Weather & Climate By Anne Marie Helmenstine, Ph.D. Chemistry Expert Ph.D., Biomedical Sciences, University of Tennessee at Knoxville B.A., Physics and Mathematics, Hastings College Dr. Helmenstine holds a Ph.D. in biomedical sciences and is a science writer, educator, and consultant. She has taught science courses at the high school, college, and graduate levels. our editorial process Facebook Facebook Twitter Twitter Anne Marie Helmenstine, Ph.D. Updated August 09, 2019 In chemistry, the law of constant composition (also known as the law of definite proportions) states that samples of a pure compound always contain the same elements in the same mass proportion. This law, together with the law of multiple proportions, is the basis for stoichiometry in chemistry. In other words, no matter how a compound is obtained or prepared, it will always contain the same elements in the same mass proportion. For example, carbon dioxide (CO2) always contains carbon and oxygen in a 3:8 mass ratio. Water (H2O) always consists of hydrogen and oxygen in a 1:9 mass ratio. Law of Constant Composition History Discovery of this law is credited to the French chemist Joseph Proust, who through a series of experiments conducted from 1798 to 1804 concluded that chemical compounds consisted of a specific composition. Considering John Dalton's atomic theory was only just beginning to explain that each element consisted of one type of atom and at the time, most scientists still believed elements could combine in any proportion, Proust's deductions were exceptional. Law of Constant Composition Example When you work with chemistry problems using this law, your goal is to look for the closest mass ratio between the elements. It's okay if the percentage is a few hundredths off. If you're using experimental data, the variation might be even greater. For example, let's say that using the law of constant composition, you want to demonstrate that two samples of cupric oxide abide by the law. Your first sample was 1.375 g cupric oxide, which was heated with hydrogen to yield 1.098 g of copper. For the second sample, 1.179 g of copper was dissolved in nitric acid to produce copper nitrate, which was subsequently burned to produce 1.476 g of cupric oxide. To work the problem, you'd need to find the mass percent of each element in each sample. It doesn't matter whether you choose to find the percentage of copper or the percentage of oxygen. You'd simply subtract one of the values from 100 to get the percent of the other element. Write down what you know: In the first sample: copper oxide = 1.375 gcopper = 1.098 goxygen = 1.375 - 1.098 = 0.277 g percent oxygen in CuO = (0.277)(100%)/1.375 = 20.15% For the second sample: copper = 1.179 gcopper oxide = 1.476 goxygen = 1.476 - 1.179 = 0.297 g percent oxygen in CuO = (0.297)(100%)/1.476 = 20.12% The samples follow the law of constant composition, allowing for significant figures and experimental error. Exceptions to the Law of Constant Composition As it turns out, there are exceptions to this rule. There are some non-stoichiometric compounds that exhibit a variable composition from one sample to another. An example is wustite, a type of iron oxide that may contain 0.83 to 0.95 iron per each oxygen. Also, because there are different isotopes of atoms, even a normal stoichiometric compound may display variations in mass composition, depending which isotope of the atoms is present. Typically, this difference is relatively small, yet it does exist and can be important. The mass proportion of heavy water as compared with regular water is an example.