Atom

I

Introduction

Atom, smallest unit of a chemical element that can exist. In ancient Greek philosophy the word “atom” was used to describe the smallest bit of matter that could be conceived of. This “fundamental particle”, to use the present-day term for this concept, was thought of as indestructible; in fact, the Greek word for atom (atomos) means “not divisible”. Knowledge about the size and nature of the atom did not begin to be acquired until long after the beginnings of experimental science in the 16th and 17th centuries. Although many of the new “experimental philosophers” believed in the reality of atoms, the progress of science owed little to the idea. The first quantitative explanation of the behaviour of matter in terms of atoms was attempted by Daniel Bernoulli in 1738, but his work was largely ignored. However, chemistry was discovering things about matter that only the idea of atoms could explain. Chemists recognized that all liquids, gases, and solids can be broken down into their ultimate components, or elements. For example, salt is a chemical compound formed when the elements sodium and chlorine react together and become joined in an intimate form known as a chemical compound. Air, by contrast, was found to consist of a mixture of the gases nitrogen and oxygen, which did not react with each other.

II

Dalton’s Theory

John Dalton, a British schoolmaster and chemist, was fascinated by the patchwork puzzle of the elements. Early in the 19th century he made studies of the way in which the various elements combine with one another to form chemical compounds. Although many other scientists, from the Greeks onward, had already speculated that the smallest units of a substance are atoms, Dalton is regarded as one of the most significant figures in atomic theory because he made the subject quantitative. He showed how these atoms link together in definite proportions. Subsequent investigations proved that atoms normally form groups called molecules. Each molecule of water, for example, consists of a single atom of oxygen and two atoms of hydrogen joined by an electrical force called a chemical bond. Water is symbolized as HOH, or as H₂O, meaning that its molecule consists of two atoms of hydrogen joined to one atom of oxygen. See Chemical Reaction.

All atoms of any given element behave in the same way chemically. Thus, from a chemical viewpoint, the atom is the smallest entity to be considered. The chemical properties of the various elements are quite different; their atoms combine in many different ways to form a multitude of different chemical compounds. Some elements, such as the noble gases helium and argon, are inert—that is, they fail to react with other elements except under special conditions. Unlike oxygen, which has a diatomic molecule (two atoms combined in a single molecule), helium and other inert gases are monatomic elements, with a single atom per molecule.

III

Avogadro’s Law

The study of gases attracted the attention of the Italian physicist Amedeo Avogadro, who in 1811 formulated an important law bearing his name. This law states that equal volumes of different gases contain the same number of molecules when compared under the same conditions of temperature and pressure. Given these conditions, two identical bottles, one filled with oxygen and the other with helium, will contain exactly the same number of molecules. Twice as many atoms of oxygen will be present, however, because oxygen is diatomic. See also Avogadro’s Law.

IV

Relative Atomic Mass

It follows from Avogadro’s law that the masses of standard volumes (that is, the densities) of different gases are proportional to the masses of the individual gas molecules. When carbon is taken as a standard and the carbon atom is assigned a value of 12.0000 atomic mass units (symbol u, formerly amu), the hydrogen atom is found to have a mass of 1.0079 u, helium 4.0026 u, fluorine 18.9984 u, and sodium 22.9898 u. Chemists sometimes quote these numbers, without mentioning any unit. These numbers were formerly called “atomic weights”, but are now called “relative atomic masses” (r.a.m.). The new term is preferred because mass is a measure of the quantity of matter in a body, which is the concept that is relevant here; weight is quite different: the force exerted on the body by the influence of gravity.

The observation that many relative atomic masses are close to whole numbers led the British chemist William Prout to suggest in 1816 that all elements might be composed of hydrogen atoms. However, subsequent measurements of relative atomic masses revealed that chlorine, for example, has a r.a.m. of 35.453 u (when carbon is taken as 12). The discovery of such fractional r.a.m. appeared to invalidate Prout’s hypothesis until a century later, when it was discovered that generally the atoms of a given element do not all have the same mass. Atoms of the same element that differ in mass are known as isotopes. In the case of chlorine two isotopes occur in nature. The atoms of one isotope (chlorine-35) have a r.a.m. close to 35 u, while those of the other (chlorine-37) have a r.a.m. close to 37 u. Experiments show that chlorine is a mixture of approximately three parts of chlorine-35 for every one part of chlorine-37. The average r.a.m. of naturally occurring chlorine is therefore approximately (3 × 35 + 37)/4 = 35.5.

During the first part of the 20th century natural oxygen was used as the standard against which atomic masses were measured; oxygen was assigned a r.a.m. of exactly 16. In the early 1960s, the international unions of chemistry and physics agreed on a new standard, assigning a r.a.m. of precisely 12 to the most commonly found isotope of carbon, carbon-12. The new standard is particularly appropriate because carbon-12 is often used as a reference standard when atomic masses are measured with mass spectrometers. The table of relative atomic masses based on carbon-12 is in close agreement with the old table based on natural oxygen.