Atom

An atom is the smallest unit of ordinary matter that retains the chemical properties of a chemical element. It consists of a dense central nucleus, made up of positively charged protons and electrically neutral neutrons, surrounded by a diffuse cloud of negatively charged electrons. When the number of protons equals the number of electrons, the atom is electrically neutral; when this balance is disturbed, the atom becomes an ion. Atoms are extraordinarily small—typically on the order of 100 picometres (10⁻¹⁰ metres) across—so small that a single drop of water contains on the order of 10²¹ of them, and their behavior can be described only through the laws of quantum mechanics rather than classical physics.

Atoms are the fundamental building blocks of chemistry and of the ordinary matter encountered in everyday life. About 118 chemical elements are recognized, each defined by the number of protons in its atoms, and every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Although atoms themselves are composed of smaller particles—quarks, electrons, and gluons—the atom marks the boundary below which the chemical identity of matter is lost: dividing an atom further yields fragments that no longer behave as the original element.

Etymology and Philosophical Origins

The English word "atom" derives from the Greek atomos (ἄτομος), meaning "uncuttable" or "indivisible," formed from the negative prefix a- and temnein, "to cut." The idea that matter is composed of discrete, indivisible units was advanced in the fifth century BCE by the Greek philosophers Leucippus and his student Democritus, who argued that all matter consists of eternal particles moving through empty space. A broadly parallel atomistic tradition developed in ancient India, notably in the Nyaya–Vaiśeṣika school. These early theories were philosophical speculations based on reasoning rather than experiment, and they competed for two millennia with Aristotle's influential doctrine that matter is continuous and composed of four elemental qualities. Despite their conceptual prescience, the ancient atomists could not test, verify, or quantify their claims, and the idea lay largely dormant until the scientific revolution.

Development of Modern Atomic Theory

The empirical foundations of atomic theory were laid during the chemical revolution of the late eighteenth century. Antoine Lavoisier's law of conservation of mass (1789) and Joseph Proust's law of definite proportions (1799) established regularities in chemical reactions that demanded explanation. In 1803 the English chemist John Dalton proposed that each chemical element consists of atoms of a single, characteristic weight, that atoms are neither created nor destroyed in chemical reactions, and that compounds form when atoms combine in simple whole-number ratios. Dalton's theory, published in A New System of Chemical Philosophy (1808), explained the law of multiple proportions and won rapid acceptance, marking the transformation of atomism from philosophy into science.

Subsequent work refined and quantified the theory. Amedeo Avogadro argued in 1811 that equal volumes of gas contain equal numbers of molecules, and the acceptance of his hypothesis—consolidated by Stanislao Cannizzaro at the Karlsruhe Congress of 1860—clarified the distinction between atoms and molecules and allowed accurate relative atomic masses to be determined. The kinetic theory of gases, developed by James Clerk Maxwell and Ludwig Boltzmann, explained the macroscopic properties of gases in terms of the motion of vast numbers of particles. The reality of atoms remained contested into the early twentieth century, but Albert Einstein's 1905 analysis of Brownian motion, confirmed experimentally by Jean Perrin, provided persuasive evidence for the discrete particulate nature of matter and yielded a reliable estimate of Avogadro's constant (approximately 6.022 × 10²³ particles per mole).

Discovery of Atomic Substructure

The closing decades of the nineteenth century revealed that atoms possess internal structure. In 1896 Henri Becquerel discovered radioactivity, subsequently studied intensively by Marie and Pierre Curie and by Ernest Rutherford, who identified alpha and beta radiation and showed that one element could spontaneously transmute into another. In 1897 J. J. Thomson demonstrated that cathode rays consist of particles far lighter than any atom—the electron—establishing for the first time that atoms are divisible. Thomson proposed the "plum pudding" model, in which electrons were embedded in a diffuse sphere of positive charge.

This model was overthrown by the Geiger–Marsden gold-foil experiment of 1909, conducted under Rutherford's direction. The unexpected deflection of some alpha particles through large angles led Rutherford to conclude in 1911 that nearly all the atom's mass and all of its positive charge are concentrated in a minute central nucleus. In 1913 Niels Bohr incorporated Max Planck's quantum hypothesis into a model in which electrons occupy discrete orbits with quantized energies, successfully explaining the spectral lines of hydrogen. In the same era, Henry Moseley's X-ray experiments showed that the atomic number—the nuclear charge, later identified with the