Law of Conservation of Mass, Definite Proportions and Dalton's Atomic Theory

Law of conservation of mass: what Lavoisier found

Purdue University dates the turning point in the study of chemical reactions to the French chemist Lavoisier’s work between 1772 and 1794. He found that the total mass of the products always equals the total mass of the starting materials that were used up. Eames, of City College of San Francisco, adds that the law carries his name, though others had used it before him.

During a reaction no matter is created or destroyed. Purdue gives the reason in terms of atoms: a reaction neither makes them nor removes them.

Practical example: the hydrogen atoms in H2 and the oxygen atoms in O2 recombine as H2O. Their number is the same before and after, and the balance shows the same mass.

Water by mass in Lavoisier and in Eames

Late in the eighteenth century Lavoisier published a list of elements. In the Traité élémentaire de chimie, from 1789, he gave water as roughly 85% oxygen and 15% hydrogen by weight. Eames gives about 89% and 11% by mass, for water made by burning hydrogen in oxygen.

Law of constant proportion (definite proportions)

At the time, Eames writes, people argued over whether elements always combine in the same ratio. Berthollet, a respected scientist, held that the constant ratios seen in experiments came from the conditions of the experiment, and that without them the ratios could vary continuously. Proust took the opposite view and eventually persuaded him. On how an observation becomes a law, see the Recap on the scientific method.

In Eames’s view Proust came out right since he tested pure compounds alone and knew how precise his own measurements were.

Practical example: in water, Eames’s percentages are about 11% hydrogen and 89% oxygen. In 100 g of water that is 11 g of hydrogen and 89 g of oxygen, about 8 g of oxygen for every gram of hydrogen (89 divided by 11, our own calculation). The law says the ratio is the same in every sample of pure water.

Eames adds that some materials, such as minerals and alloys, can have variable proportions. The law of definite proportions applies to pure compounds.

Law of multiple proportions

Dalton went through the data of Proust, Lavoisier and others and spotted a pattern. Eames explains that the law of multiple proportions compares two or more compounds of the same elements in different proportions, and shows that the elements combine in small whole-number ratios.

Practical example: carbon combines with oxygen in mass ratios of 3 to 4 and 3 to 8. So 3 g of carbon bind 4 g or 8 g of oxygen, and one compound holds exactly twice the oxygen of the other. Nitrogen and oxygen give 7 to 4, 7 to 8 and 7 to 16, with the oxygen doubling at each step.

The three laws side by side

Law	What it states	Example from the sources
Conservation of mass (Lavoisier)	The total mass of the products equals that of the reactants	Hydrogen and oxygen recombining as water
Definite proportions (Proust)	In a pure compound the mass proportions between the elements do not change	Water: 11% hydrogen and 89% oxygen by mass
Multiple proportions (Dalton)	Two elements forming several compounds combine in small whole-number ratios	Carbon and oxygen: 3 to 4 and 3 to 8

Dalton’s atomic theory and the billiard ball model

For Eames, the three laws made sense if every element is made of atoms that combine as whole atoms. That is the idea behind Dalton’s atomic theory, also called the Dalton model: matter is made of atoms, which are not created, not destroyed and do not change in chemical reactions. All atoms of one element are identical, and the atoms of each element have a mass of their own, different from that of the other elements.

The AIP gives a five-point version, taken from Partington (1962). One of the five points says the particle of a compound holds a fixed number of atoms of its elements; Partington’s list goes on with six more statements, so the five points are only part of it. Eames compares the atom to a ball, small and distinct from the others, and the nickname billiard ball model echoes that image.

John Dalton (1766-1844) was born in Cumberland, England, into a modest Quaker family, and spent nearly all his life as a teacher and public lecturer. The Science History Institute tells us he came to atomism through meteorology. The paper with the first form of the theory, titled On the Absorption of Gases by Water and Other Liquids, was read on 21 October 1803. According to the AIP, William Henry and Thomas Thomson deserve credit for the first accounts, before Dalton published it himself in 1808, and the discovery was not free of controversy: Dalton was accused of using earlier work without citing it. The same source says it is far from obvious how Dalton arrived at the idea of atoms combining in small whole numbers.

Dalton gathered his theories in A New System of Chemical Philosophy, published between 1808 and 1827. He worked out atomic weights from the percentage composition of compounds, with an arbitrary system to decide the likely structure of each. His rule: for two elements, A and B, that combine, the first compound has one atom of each, and the next has one atom of A and two of B.

Practical example: with that system Dalton gave water the formula HO, which is wrong compared with the modern H2O.

What was corrected in Dalton’s model

According to Eames, the atoms of an element are neither all identical nor indivisible: they are made of three main kinds of particle, electrons, protons and neutrons. Protons and neutrons together form a tiny, dense center, the nucleus. In 1897 J. J. Thomson suggested that cathode rays were particles with a much larger charge or a much smaller mass than hydrogen ions, and many were skeptical. Electrons turned out to be part of every kind of atom. The evidence against Dalton’s identical, indivisible atoms grew strong.

An element is defined by its number of protons, the atomic number (see also the periodic table). Isotopes are atoms with the same number of protons and a different number of neutrons. Chemistry follows the atomic number closely, so isotopes behave almost alike in reactions, yet their masses differ.

Practical example: Aston found that neon atoms have mass 20 or 22, not 20.2, the value accepted at the time. They are two versions of one element: same chemical properties, different masses.

Dalton’s postulate	What Eames reports
Atoms are indivisible	They have smaller parts: electrons, protons, neutrons
Atoms of an element are identical	Isotopes have different masses