Chemical element

A chemical element is a pure chemical substance consisting of a single type of atom distinguished by its atomic number, which is the number of protons in its atomic nucleus. All ordinary matter in the universe is composed of chemical elements, either in their native form or combined into chemical compounds. As of 2025, a total of 118 elements have been officially recognized, ranging from hydrogen (atomic number 1) to oganesson (118), with the first 94 occurring naturally on Earth in appreciable quantities and the remainder being synthetic products of nuclear reactions. The concept of the element is foundational to chemistry, physics, and materials science, providing the organizing principle behind the periodic table and the systematic understanding of the properties and interactions of matter.

History of the concept

The idea that matter is composed of discrete, fundamental constituents dates back to ancient philosophies. Pre‑Socratic Greek thinkers such as Empedocles proposed that all matter was a combination of four “roots” — earth, air, fire, and water — while Democritus and Leucippus advanced an atomic theory positing indivisible particles called atomos. These early notions, however, were speculative and lacked empirical foundation.

In the 17th century, Robert Boyle challenged the classical four‑element theory in The Sceptical Chymist (1661), arguing that an element should be defined experimentally as a substance that cannot be broken down into simpler substances by chemical means. This operational definition guided chemical research for the next two centuries. Antoine Lavoisier, in his Traité Élémentaire de Chimie (1789), published a list of 33 simple substances that he regarded as elements, including light and caloric, which were later discarded. Importantly, Lavoisier’s emphasis on quantitative measurement and the conservation of mass laid the groundwork for a modern chemical ontology.

John Dalton’s atomic theory (1808) provided the first theoretical framework linking the concept of an element to a specific type of atom with a characteristic atomic weight. Throughout the 19th century, the identification of new elements accelerated, and in 1869 Dmitri Mendeleev published his periodic table, arranging the known elements by increasing atomic weight and grouping them according to recurring chemical properties. Mendeleev’s table predicted the existence and properties of several then‑unknown elements, cementing the periodic law as a central principle of chemistry. The discovery of the electron by J. J. Thomson (1897), the nuclear model of the atom by Ernest Rutherford (1911), and the determination of atomic numbers by Henry Moseley (1913) finally established that the defining property of an element is the nuclear charge, rather than atomic weight, resolving inconsistencies in the periodic order.

Definition and atomic structure

The modern definition of a chemical element is grounded in atomic physics. Every atom consists of a dense nucleus containing positively charged protons and electrically neutral neutrons, surrounded by a cloud of negatively charged electrons. The number of protons in the nucleus, termed the atomic number (denoted Z), uniquely identifies the element. For example, all atoms with one proton are hydrogen, all with six protons are carbon, and all with 92 protons are uranium. While atoms of the same element always have the same number of protons, they can vary in the number of neutrons, giving rise to different isotopes of that element. The sum of protons and neutrons is the mass number (A). The chemical behavior of an element is largely determined by its electron configuration, which in turn depends on the atomic number and governs bonding, reactivity, and placement in the periodic table.

An element in its standard state may exist as individual atoms (as in the noble gases), as molecules composed of multiple atoms of the same element (such as diatomic oxygen, O₂, or octasulfur, S₈), or as extended networks (such as metallic lattices or covalent network solids like diamond). The term “element” refers to the species of atoms defined by Z, irrespective of the physical aggregation state of the substance.

Properties

General physical and chemical properties

Each element exhibits a characteristic set of physical and chemical properties, including electronegativity, ionization energy, atomic radius, melting point, boiling point, density, and common oxidation states. These properties show periodic trends across the periodic table — for instance, ionization energy generally increases across a period and decreases down a group. The electron configuration determines the element’s valency and the types of chemical bonds it can form, which in turn dictate the stoichiometry and structure of its compounds.

Elements may be gases, liquids, or solids at room temperature. Most are metallic, possessing high electrical and thermal conductivity, malleability, ductility, and a lustrous appearance. Nonmetals display a wide range of contrasting behaviors: some are gases at ambient conditions, others are brittle solids, and they are generally poor conductors of heat and electricity. A small number of elements, such as silicon and germanium, exhibit semiconducting properties intermediate between metals and nonmetals.

Isotopes and atomic mass

Most elements occur naturally as a mixture of several isotopes. The atomic weight (or standard atomic weight) listed for an element in the periodic table is a weighted average of the masses of its naturally occurring isotopes, taking into account their relative abundances. For synthetic elements or elements with no stable isotopes, the mass number of the most stable known isotope is often cited in brackets. The existence of isotopes explains why the atomic weights of some elements (e.g., chlorine, approximately 35.45) are not close to integer multiples of the hydrogen mass.

Isotopes are critical in many fields. Radioactive isotopes, or radioisotopes, decay by emitting radiation and are used in medicine (diagnostic imaging, cancer therapy), archaeology (radiocarbon dating), and energy production (nuclear fuel). Stable isotopes serve as tracers in environmental and biological studies.

Nuclear stability

The stability of an atomic nucleus depends on the balance between protons and neutrons. Elements with atomic numbers up to 82 (lead) have at least one stable isotope, with the exceptions of technetium (43) and promethium (61), which have only radioactive isotopes. Bismuth (83) was long thought to be stable, but its only naturally occurring isotope was found in 2003 to undergo alpha decay with an extremely long half‑life. All elements beyond lead (Z > 82) are exclusively radioactive. The transuranium elements (Z > 92) do not occur naturally except in trace amounts as decay products of uranium and thorium and are predominantly synthesized in laboratories.

Classification

Metals, nonmetals, and metalloids

The most fundamental classification divides elements into metals, nonmetals, and metalloids (or semimetals). Metals, which constitute the majority of elements, are characterized by a tendency to lose electrons and form cations, metallic bonding in bulk, and the typical physical properties mentioned earlier. Nonmetals tend to gain or share electrons, forming anions or covalent compounds, and encompass elements essential to organic chemistry, such as carbon, nitrogen, and oxygen. Metalloids (e.g., boron, silicon, germanium, arsenic, antimony, tellurium) exhibit properties intermediate between metals and nonmetals and are important in semiconductor technology.

Blocks and groups of the periodic table

The periodic table arranges elements in order of increasing atomic number into rows (periods) and columns (groups) such that elements with similar chemical behaviors fall into the same group. The table can be divided into blocks based on the electron subshell being filled:

• s‑block: Groups 1 (alkali metals) and 2 (alkaline earth metals), plus helium. These elements have their outermost electrons in an s orbital and are highly reactive metals (except helium).
• p‑block: Groups 13 to 18, incorporating metals, metalloids, and nonmetals, including the boron, carbon, nitrogen, chalcogen, halogen, and noble gas groups. The halogens (group 17) are particularly reactive nonmetals, while the noble gases (group 18) are largely inert.
• d‑block: Groups 3 to 12, the transition metals, which have partially filled d orbitals. They typically form colored compounds, exhibit variable oxidation states, and are often good catalysts.
• f‑block: The lanthanides (elements 57–71) and actinides (89–103), often displayed as two rows separate from the main table. The lanthanides are reactive metals with similar chemical properties, while the actinides are all radioactive and include the naturally occurring uranium and thorium alongside synthetic elements like plutonium and americium.

Other groupings

Additional groupings are used in specific contexts, such as the platinum group metals, precious metals, refractory metals, coinage metals, and rare earth elements (which include scandium, yttrium, and the lanthanides). Biochemists frequently refer to the CHNOPS elements (carbon, hydrogen, nitrogen, oxygen, phosphorus, sulfur) as the primary building blocks of life.

Occurrence and abundance

Hydrogen is by far the most abundant element in the universe, accounting for roughly 74% of all baryonic mass, followed by helium (about 24%). All heavier elements together constitute less than 2%, having been produced primarily by stellar nucleosynthesis and supernova explosions. Nuclear fusion in stars fuses hydrogen into helium, and successive fusion processes build elements up to iron and nickel. Elements heavier than iron are predominantly formed by neutron‑capture processes (s‑process and r‑process) in stars and supernovae.

On Earth, the composition differs dramatically. The Earth’s core is largely iron and nickel. The crust is dominated by oxygen (about 46% by mass), silicon, aluminum, iron, calcium, sodium, potassium, and magnesium, mostly in the form of silicate minerals. The atmosphere is approximately 78% nitrogen and 21% oxygen, with trace amounts of argon, carbon dioxide, and other gases. Seawater contains dissolved ions of sodium, chlorine, magnesium, sulfur, calcium, and potassium, among others. The human body is composed mainly of oxygen, carbon, hydrogen, nitrogen, calcium, and phosphorus.

A few elements occur natively in uncombined form, including gold, platinum, silver, copper, sulfur, carbon (as graphite and diamond), and the noble gases. Most elements, however, are found in chemical combination as ores, from which they must be extracted by physical or chemical processes.

Discovery of the elements

The timeline of element discovery spans from ancient times to the 21st century. Carbon, sulfur, iron, copper, silver, gold, tin, antimony, mercury, and lead were known to ancient civilizations and used without recognition of their elemental nature. Phosphorus was the first element to be chemically discovered in modern times (Hennig Brand, 1669). During the 18th and 19th centuries, the use of electrolysis, flame spectroscopy, and systematic chemical analysis led to the isolation of numerous metals, including potassium, sodium, calcium, magnesium, and many others. Spectroscopy enabled the discovery of elements by their characteristic emission lines; helium, for instance, was first detected in the solar spectrum before being found on Earth. The discovery of radioactivity by Henri Becquerel and the subsequent isolation of polonium and radium by Marie and Pierre Curie opened the door to the identification of many unstable elements. The transuranium elements were synthesized from 1940 onward, starting with neptunium and plutonium, and the most recent additions to the periodic table — nihonium, moscovium, tennessine, and oganesson — completed the seventh period and were officially named in 2016.

Synthetic elements

Elements with atomic numbers above 94 (plutonium) are exclusively synthetic, produced in nuclear reactors or particle accelerators through nuclear fusion, neutron capture, or heavy‑ion bombardment. Their nuclei are highly unstable, with half‑lives ranging from microseconds to millions of years. The creation and identification of new elements is an ongoing research endeavor, with experiments aimed at producing elements in the hypothetical “island of stability” around Z ≈ 120–126, where nuclear shell effects may confer greater longevity. All synthetic elements to date have been produced in quantities too small for bulk chemical analysis, so their chemistry is often studied on an atom‑at‑a‑time basis using rapid radiochemical techniques.

List of elements

The 118 confirmed elements are each assigned a unique name, chemical symbol, and atomic number. Hydrogen (H, 1) through oganesson (Og, 118) are all formally recognized by the International Union of Pure and Applied Chemistry (IUPAC), which is the authority responsible for naming new elements and setting criteria for their discovery. The names often honor mythological figures, minerals, places, or scientists (e.g., curium, einsteinium, californium). The symbols, usually one or two letters, are derived from the element’s English or Latin name (e.g., Na for sodium from natrium, Fe for iron from ferrum). A full tabulation of the elements is traditionally presented as the periodic table, which remains one of the most powerful and concise summaries of chemical knowledge, illustrating the relationships and periodicity of elemental properties.