Chemical element

A chemical element is a fundamental chemical substance composed of atoms that all share the same number of protons in their nuclei. This proton count, known as the atomic number, uniquely identifies each element and governs its position in the periodic table, its electronic structure, and its characteristic chemical behaviour. Elements form the basic building blocks of all matter, combining in fixed or variable ratios to create molecules, compounds, and a vast range of natural and synthetic materials.

Definition and Fundamental Characteristics

An element is defined by its atoms having a constant atomic number. For example, all oxygen atoms contain eight protons, giving oxygen its atomic number of 8. While atoms of a given element have identical numbers of protons, they may differ in the number of neutrons contained in their nuclei. These variants are known as isotopes. Although isotopes of an element can have different masses, their chemical behaviour is usually indistinguishable because they possess the same electronic configuration.
Elements may occur as individual atoms, diatomic molecules, or larger homonuclear clusters. Hydrogen, for instance, naturally forms diatomic molecules (H₂). In contrast, chemical compounds consist of atoms of different elements combined in fixed proportions, whereas mixtures contain different substances that are not chemically bonded, such as the gaseous combination of nitrogen and oxygen in the Earth’s atmosphere.
Historically, the term chemical element denoted a substance that could not be decomposed by chemical means. This view, although simplified, remains useful in distinguishing elements from compounds. The development of atomic theory and the understanding of isotopes refined the definition in the twentieth century, leading to the modern distinction between elements as types of atoms and elementary substances consisting of those atoms in pure form.

Natural Occurrence and Distribution of Elements

Nearly all baryonic matter in the universe is composed of chemical elements. Exceptions include exotic astronomical objects, such as neutron stars, which consist largely of subatomic particles. Hydrogen and helium dominate the cosmos, formed primarily by Big Bang nucleosynthesis. Heavier elements were produced by stellar processes, including fusion in stars, supernovae, and neutron-star collisions.
On Earth, only a small number of elements occur naturally in uncombined form. Gold and silver are classic examples, found as native metals. Most elements, however, occur in compounds or mixtures. For instance, the atmosphere is chiefly composed of molecular nitrogen and oxygen, alongside water vapour, carbon dioxide, and argon — a chemically inert noble gas.
By late 2016, the International Union of Pure and Applied Chemistry (IUPAC) had formally recognised 118 elements. Ninety-four occur naturally, while the remainder are synthetic elements produced in particle accelerators or nuclear reactors. Naturally occurring elements with atomic numbers greater than 82 exhibit varying degrees of radioactivity. Some, such as bismuth, thorium, and uranium, possess isotopes with extremely long half-lives, allowing them to survive from the formation of the Solar System.
Beyond plutonium (atomic number 94), all elements are highly unstable and found only as artificially produced isotopes with short half-lives. Research continues to extend the periodic table, synthesise superheavy elements, and explore the theoretical limits of nuclear stability.

The Periodic Table and Systematic Organisation

The periodic table arranges elements in order of increasing atomic number, grouping them according to recurring patterns of physical and chemical properties. This organisation reveals relationships among families of elements, such as the alkali metals, halogens, noble gases, and transition metals. Dmitri Mendeleev’s pioneering work in 1869 established the framework for predicting undiscovered elements and understanding the periodicity of elemental behaviour.
Modern periodic tables incorporate electron configuration and quantum theory, explaining trends such as atomic radius, ionisation energy, electronegativity, and bond formation. These patterns allow chemists to anticipate the properties of new compounds and estimate the stability of hypothetical elements.

Atomic Number and Chemical Behaviour

The atomic number, symbolised as Z, defines an element more fundamentally than its mass. Each element’s chemical properties are governed by its electrons, which are arranged in atomic orbitals determined by the positive charge of the nucleus. Although isotopes differ in mass due to varying neutron numbers, they share the same electron configuration and therefore exhibit nearly identical chemistry. Hydrogen is the major exception, as the substantial mass differences between protium, deuterium, and tritium yield observable kinetic isotope effects.
Because atomic number determines electron structure, it remains the defining criterion for identifying and classifying elements. The mass number, by contrast, reflects an individual isotope rather than the element as a whole.

Isotopes and Nuclear Properties

Isotopes are varieties of an element whose atoms contain the same number of protons but differing numbers of neutrons. Carbon offers a typical example, with isotopes containing six, seven, or eight neutrons, known respectively as carbon-12, carbon-13, and carbon-14. Natural carbon is predominantly carbon-12, with smaller proportions of carbon-13 and minute traces of carbon-14.
Stable isotopes are those for which no radioactive decay has been detected. A majority of naturally occurring elements possess multiple stable isotopes. Radioactive isotopes — or radioisotopes — spontaneously transform into other nuclides via decay processes such as alpha decay, beta decay, inverse beta decay, or, in heavy elements, spontaneous fission. These nuclear transformations alter the atomic number and thereby change one element into another.
Many radioisotopes, especially of heavy elements, do not occur naturally and are produced artificially for applications ranging from medical imaging to industrial tracing and nuclear research.

Discovery, Synthesis, and Ongoing Expansion

The identification of elements has proceeded gradually from antiquity to modern times. Early societies recognised elemental substances such as sulphur, copper, carbon, and gold, though without the concept of atomic theory. Alchemical traditions attempted classifications that anticipated modern ideas but lacked consistency.
Substantial progress occurred during the eighteenth and nineteenth centuries, leading to the discovery of gases, metals, and non-metals through systematic experimentation. Mendeleev’s periodic table accelerated this process, enabling predictions of undiscovered elements’ properties.
Twentieth- and twenty-first-century research has greatly expanded the periodic table through nuclear synthesis. Elements such as technetium and promethium, at first thought not to occur naturally, were later identified in trace amounts. Superheavy elements, including tennessine and oganesson, have been synthesised only in specialised laboratories, sometimes from just a few observed decay events.