It states that when elements combine they do so in a ratio of small whole numbers. For example, carbon and oxygen react to form CO or CO2, but not CO1.3 for instance. Furthermore, it states that if two elements form more than one compound between them then the ratios of the masses of the second element combined with a fixed mass of the first element will also be in ratios of small whole numbers.
English chemist John Dalton first expressed this observation in 1803 and it is sometimes called Dalton's Law, though that more usually refers to his law of partial pressures. A few years previously, the French chemist Joseph Proust had proposed the law of definite proportions, which expressed that the elements combined to form compounds in certain well-defined proportions, rather than mixing in just any proportion. Careful study of the actual numerical values of these proportions led Dalton to propose his law of multiple proportions. This was an important step toward the atomic theory that he would propose later that year, and it laid the basis for the chemical formulas for compounds.
There are over 18 million known substances in our world. We will begin by assuming that all materials are made from elements, materials which cannot be decomposed into simpler substances. We will assume that we have identified all of these elements, and that there is a very small number of them. All other pure substances, which we call compounds, are made up from these elements and can be decomposed into these elements. For example, metallic iron and gaseous oxygen are both elements and cannot be reduced into simpler substances, but iron rust, or ferrous oxide, is a compound which can be reduced to elemental iron and oxygen. The elements are not transmutable: one element cannot be converted into another. Finally, we will assume that we have demonstrated the Law of Conservation of Mass.
We could, of course, jump directly to the answers to these questions by stating that the elements themselves are composed of atoms: indivisible, identical particles distinctive of that element. Then a compound is formed by combining the atoms of the composite elements. Certainly, the Law of Conservation of Mass would be easily explained by the existence of immutable atoms of fixed mass. However, if we do decide to jump to conclusions and assume the existence of atoms without further evidence (as did the leading chemists of the seventeenth and eighteenth centuries), it does not lead us anywhere. What happens to iron when, after prolonged heating in air, it converts to iron rust? Why is it that the resultant combination of iron and air does not maintain the properties of either, as we would expect if the atoms of each are mixed together? An atomic view of nature would not yet provide any understanding of how the air and the iron have interacted or combined to form the new compound, and we can't make any predictions about how much iron will produce how much iron rust. There is no basis for making any statements about the properties of these atoms. We need further observations.
Second we examine ammonia, which is a combination of nitrogen and hydrogen with the mass proportion of 7 : 1.5 nitrogen to hydrogen. If this is one nitrogen combined with one hydrogen, we would expect that a nitrogen atom mass is 4.67 times that of a hydrogen atom mass. These two expectations predict a relationship between the mass of an oxygen atom and the mass of a hydrogen atom. If the mass of an oxygen atom is 1.14 times the mass of a nitrogen atom and if the mass of a nitrogen atom is 4.67 times the mass of a hydrogen atom, then we must conclude that an oxygen atom has a mass which is 1.14 × 4.67 = 5.34 times that of a hydrogen atom. But there is a problem with this calculation. The third line of Table 1 shows that the compound formed from hydrogen and oxygen is water, which is found to have mass proportion 8:1 oxygen to hydrogen.
Our expectation should then be that an oxygen atom mass is 8.0 times a hydrogen atom mass. Thus the three measurements in Table 1 appear to lead to contradictory expectations of atomic mass ratios. How are we to reconcile these results? Mass Relationships for Hydrogen, Nitrogen, Oxygen Compounds Compound Total Mass Mass of Hydrogen Mass of Nitrogen Mass of Oxygen "Expected" Relative Atomic Mass of Hydrogen "Expected" Relative Atomic Mass of Nitrogen "Expected" Relative Atomic Mass of Oxygen Nitric Oxide 15.0 g - 7.0 g 8.0 g - 7.0 8.0 Ammonia 8.5 g 1.5 g 7.0 g - 1.5 7.0 - Water 9.0 g 1.0 g - 8.0 g 1.0 - 8.0 One possibility is that we were mistaken in assuming that there are atoms of the elements which combine to form the different compounds. If so, then we would not be surprised to see variations in relative masses of materials which combine. Another possibility is that we have erred in our reasoning. Looking back, we see that we have to assume how many atoms of each type are contained in each compound to find the relative masses of the atoms. In each of the above examples, we assumed the ratio of atoms to be 1:1 in each compound. If there are atoms of the elements, then this assumption must be wrong, since it gives relative atomic masses which differ from compound to compound. How could we find the correct atomic ratios? It would help if we knew the ratio of the atomic masses: for example, if we knew that the oxygen to hydrogen mass ratio were 8:1, then we could conclude that the atomic ratio in water would be 1 oxygen and 1 hydrogen. Our reasoning seems to circular: to know the atomic masses, we must know the formula of the compound (the numbers of atoms of each type), but to know the formula we must know the masses. Which of these possibilities is correct? Without further observations, we cannot say for certain whether matter is composed of atoms or not.
The masses of oxygen appearing in these compounds are in simple whole number ratios when we take a fixed amount of nitrogen. The appearance of these simple whole numbers is very significant. These integers imply that the compounds contain a multiple of a fixed unit of mass of oxygen. The simplest explanation for this fixed unit of mass is that oxygen is particulate. We call the fixed unit of mass an atom. We now assume that the compounds have been formed from combinations of atoms with fixed masses, and that different compounds have differing numbers of atoms. The mass ratios make it clear that oxide B contains twice as many oxygen atoms (per nitrogen atom) as does oxide C and half as many oxygen atoms (per nitrogen atom) as does oxide A. The simple mass ratios must be the result of the simple ratios in which atoms combine into molecules. If, for example, oxide C has the molecular formula NO, then oxide B has the formula NO2, and oxide A has the formula NO4. There are other possibilities: if oxide B has molecular formula NO, then oxide A has formula NO2, and oxide C has formula N2O. Or if oxide A has formula NO, then oxide B has formula N2O and oxide C has formula N4O. These three possibilities are listed in Table 2. Possible Molecular Formula for Nitrogen Oxides Assuming that: Oxide C is NO Oxide B is NO Oxide A is NO Oxide A is NO4 NO2 NO Oxide B is NO2 NO N2O Oxide C is NO N2O N4O We don't have a way (from these data) to know which of these sets of molecular formula are right. But we can assert that either one of them or one analogous to them is right. Similar data are found for any set of compounds formed from common elements. For example, there are two oxides of carbon, one with oxygen to carbon mass ratio 1.33:1 and the other with mass ratio 2.66:1. The second oxide must have twice as many oxygen atoms, per carbon atom, as does the first. The general statement of this observation is the Law of Multiple Proportions.
· the elements are composed of identical atoms · all atoms of a single element have the same characteristic mass · these number and masses of these atoms do not change during a chemical transformation · compounds consist of identical molecules formed of atoms combined in simple whole number ratios
In short, the law of multiple proportions states: "when 2 elements form more than 1 compound, the different masses of 1 element that are combined with the same mass of the other element are in a ratio of small whole numbers".