An atom is the smallest unit of matter that keeps the identity of a chemical element. Every atom has a tiny, dense, positively charged nucleus made of protons and neutrons. A cloud of negatively charged electrons surrounds the nucleus. Atoms are the building blocks of ordinary solids, liquids, gases and plasmas. Their interactions explain much of chemistry and a large part of physics. A typical atom is about 0.1 nanometres across. The nucleus is roughly 100,000 times smaller than that but holds more than 99.9% of the atom's mass. The word comes from the Greek atomos, meaning "uncuttable." Atoms can in fact be divided, but a chemical process cannot break them apart.
Early ideas
The idea that matter is made of tiny, indivisible particles goes back to ancient Greece. In the 5th century BCE, Leucippus and his student Democritus argued that the world consists of countless atoms moving through empty space. In their view, differences in shape, arrangement and position explained the variety of things. Thinkers of the Vaiśeṣika school in ancient India developed similar ideas. Aristotle rejected atomism in favor of continuous matter made of four elements, and his view dominated European thought for centuries. Atomism stayed largely philosophical until the rise of modern science. Seventeenth-century "corpuscular" theories, such as those of Pierre Gassendi and Robert Boyle, brought it back.
Atomic theory in chemistry
The atom became a scientific concept through quantitative chemistry. Antoine Lavoisier's careful measurement of masses in chemical reactions and Joseph Proust's law of definite proportions prepared the way. Around 1803, John Dalton proposed the following:
- Each element consists of identical atoms with a characteristic mass.
- Atoms of different elements differ from one another.
- Compounds form when atoms combine in fixed whole-number ratios.
- Chemical reactions rearrange atoms but do not create or destroy them.
His theory explained the law of multiple proportions. For example, the masses of oxygen that combine with a fixed mass of carbon in carbon monoxide and carbon dioxide are in a 1:2 ratio. Amedeo Avogadro distinguished atoms from molecules in 1811. Later, Dmitri Mendeleev's periodic table of 1869 arranged the elements by recurring chemical properties.
Some 19th-century scientists still treated atoms as a convenient fiction. The question was settled after Albert Einstein explained Brownian motion in 1905 as the result of molecular collisions, and Jean Perrin confirmed his predictions experimentally in 1908.
Discovery of internal structure
Around 1900, experiments showed that atoms are not indivisible after all. In 1897 J. J. Thomson identified the electron as a particle far lighter than any atom. He proposed the "plum pudding" model, in which electrons sit inside a diffuse sphere of positive charge. The discovery of radioactivity in 1896 showed that atoms can change spontaneously into other kinds of atoms.
In 1909, Hans Geiger and Ernest Marsden fired alpha particles at thin gold foil, working under Ernest Rutherford at the University of Manchester. Most of the particles passed straight through, but a few bounced back sharply. In 1911 Rutherford concluded that almost all of an atom's mass and all of its positive charge sit in a tiny central nucleus. In 1913 Henry Moseley showed that an element's position in the periodic table depends on the charge of its nucleus. In 1932 James Chadwick discovered the neutron.
Quantum models
Rutherford's model could not explain why atoms are stable or why they emit light only at certain wavelengths. Under classical electromagnetism, an electron orbiting a nucleus should radiate energy and spiral inward. In 1913 Niels Bohr proposed that electrons occupy only certain allowed orbits with fixed energies. In his model, an atom emits or absorbs light only when an electron jumps between these levels. This reproduced the spectrum of hydrogen.
Louis de Broglie's 1924 idea that particles behave as waves led to quantum mechanics, and Erwin Schrödinger applied the new theory to the atom in 1926. In the modern picture, electrons do not follow definite paths. They are described by orbitals, which give the probability of finding an electron in a given region. Four quantum numbers label each electron's state. The Pauli exclusion principle says no two electrons in an atom can share the same set of quantum numbers. Together, these rules explain how electron shells fill and why the periodic table has the shape it does.
Structure and properties
An atom's atomic number (Z) is the number of protons in its nucleus, and this number defines the element. Its mass number is the total number of protons and neutrons. Atoms of the same element with different numbers of neutrons are called isotopes. Some isotopes are stable, while others are radioactive.
A proton has about 1,836 times the mass of an electron, and a neutron is slightly heavier than a proton. Protons and neutrons are not fundamental particles. Each is made of three quarks bound together by the strong interaction, which also holds the nucleus together against the electrical repulsion between protons. Electrons are fundamental particles in the Standard Model of particle physics.
A neutral atom has the same number of electrons as protons. An atom that gains or loses electrons becomes an ion. The outermost electrons, called valence electrons, decide how an atom forms chemical bonds with other atoms. Because the nucleus is so small, most of an atom's volume is taken up by its electron cloud.
As of the 2020s, 118 elements have been identified, from hydrogen (Z = 1) to oganesson (Z = 118). Elements heavier than uranium are mostly made artificially and are usually very short-lived.
Nuclear processes and observation
Chemical reactions involve only the electrons, but nuclear reactions change the nucleus itself. These include radioactive decay, nuclear fission (discovered in 1938) and nuclear fusion. They release far more energy per atom than chemical reactions do, and they are studied in nuclear physics.
Atoms are much smaller than the wavelengths of visible light, so they cannot be seen with an ordinary optical microscope. Since the 1980s, however, the scanning tunneling microscope and related instruments have imaged individual atoms and even moved them one at a time. Today, precise measurements of atomic transitions are the basis of atomic clocks and of the SI definition of the second.
References
- The History of the Atomic Modelalloprof.qc.ca
- Bohr's Atomic Modelflexbooks.ck12.org
- History of the Atomic Model: Democritus to Quantumthecalculatedchemist.com
- The Bohr model: The famous but flawed depiction of an atomspace.com
- 1.14: The Nuclear Atom - Chemistry LibreTextschem.libretexts.org
- New Types of Hydrogenlike matter Composed of Electron(s) and Meson(s)arxiv.org