Laws of Chemical Combination and Dalton's Atomic Theory

By the end of the eighteenth century, careful weighing experiments on chemical reactions had revealed several

striking regularities. These are known as the laws of chemical combination, and together they became the

experimental foundation on which John Dalton built the first atomic theory of matter.

1. Law of Conservation of Mass (Antoine Lavoisier, 1789): In a chemical reaction, matter can neither be created nor destroyed — the total mass of the reactants exactly equals the total mass of the products, when

the reaction is carried out in a closed system. If $56\ \text{g}$ of calcium oxide and $44\ \text{g}$ of carbon

dioxide combine to form calcium carbonate, exactly $100\ \text{g}$ of calcium carbonate is formed — no mass is

lost or gained. This law can only be verified when the reaction is carried out in a sealed vessel; if a gas is

allowed to escape (as when a metal carbonate is heated in an open dish and $\text{CO}_2$ bubbles away), the

solid residue alone will appear lighter, but the total mass of everything produced, including the escaped

gas, is unchanged.

2. Law of Definite Proportions (Constant Composition) (Joseph Proust, 1799): A given chemical compound always contains exactly the same elements combined together in the same fixed proportion by mass, regardless of the source of the compound or how it was prepared. Water obtained from a river, from melted ice, or made

by burning hydrogen in oxygen in a laboratory, always contains hydrogen and oxygen in the same mass ratio,

$1 : 8$. This law tells us that a compound is not an arbitrary blend — its composition is a fixed chemical

identity.

3. Law of Multiple Proportions (John Dalton, 1803): When two elements combine to form more than one compound, the different masses of one element that combine with a fixed mass of the other element are in a ratio of small whole numbers. Carbon and oxygen form two well-known oxides: carbon monoxide, $\text{CO}$, and

carbon dioxide, $\text{CO}_2$. In $\text{CO}$, $12\ \text{g}$ of carbon combines with $16\ \text{g}$ of oxygen;

in $\text{CO}_2$, the same $12\ \text{g}$ of carbon combines with $32\ \text{g}$ of oxygen. The two masses of

oxygen, $16\ \text{g}$ and $32\ \text{g}$, are in the simple ratio $1:2$ — exactly what this law predicts.

4. Gay-Lussac's Law of Gaseous Volumes (Joseph Gay-Lussac, 1808): When gases react with one another, the volumes of the reacting gases and the volumes of the gaseous products, measured at the same temperature and pressure, are in a ratio of small whole numbers. For example, in the formation of ammonia,

$\text{N}_2(g) + 3\text{H}_2(g) \rightarrow 2\text{NH}_3(g)$, one volume of nitrogen always combines with three

volumes of hydrogen to give two volumes of ammonia — the volume ratio $1:3:2$ is fixed, no matter how much gas

is actually used.

5. Avogadro's Law (Amedeo Avogadro, 1811): Equal volumes of all gases, at the same temperature and pressure, contain equal numbers of molecules. This was the crucial insight that finally explained

Gay-Lussac's law in terms of particles rather than just volumes, and it is the law that ultimately connects gas

volumes to the mole concept (Section 1.4): at standard temperature and pressure (STP), one mole of any ideal

gas occupies the same volume, $22.4\ \text{L}$.

Dalton's Atomic Theory (1808). To explain these laws, John Dalton proposed that matter is made of

indivisible particles called atoms. His main postulates were:

1. All matter is made up of tiny, indivisible particles called atoms.
2. Atoms of the same element are identical in mass and properties; atoms of different elements differ in mass and properties.
3. Atoms can neither be created nor destroyed in a chemical reaction (this immediately explains the law of conservation of mass — if atoms are only rearranged, not created or destroyed, the total mass before and after a reaction must be the same). …