List of chemical elements

The list of chemical elements is a systematic compilation of the 118 recognized chemical elements, each defined as a pure substance consisting of atoms with the same atomic number—the number of protons in their nuclei.^[1] These elements, from hydrogen (atomic number 1) to oganesson (atomic number 118), serve as the fundamental building blocks of all matter, with their properties determined by atomic structure and electron configuration.^[2] The list is conventionally presented in the periodic table, which arranges elements in rows (periods) and columns (groups) to highlight recurring trends in reactivity, electronegativity, and other characteristics.^[2] The discovery and recognition of chemical elements have evolved over millennia, starting with ancient observations of naturally occurring metals like gold, silver, copper, iron, tin, lead, mercury, and sulfur, known since prehistoric times for their utility in tools and artifacts.^[3] By the 18th and 19th centuries, systematic chemical analysis led to the identification of additional elements such as oxygen (1774), hydrogen (1766), and nitrogen (1772), culminating in Dmitri Mendeleev's 1869 periodic table that predicted undiscovered elements based on gaps in atomic weights.^[4] Today, 94 elements occur naturally on Earth—primarily through primordial formation in stars or cosmic events, with some trace amounts from radioactive decay—while the remaining 24 transuranic elements (atomic numbers 95–118) are synthetic, created via nuclear reactions in laboratories since the 1940s.^[5][6] The International Union of Pure and Applied Chemistry (IUPAC), in collaboration with the International Union of Pure and Applied Physics (IUPAP), oversees the verification of new element discoveries, assigns official names (often honoring scientists, places, or mythological figures), and updates standard atomic weights to reflect isotopic variations in natural samples.^[2] This authoritative process ensures the list remains current, with the most recent additions—nihonium (113), moscovium (115), tennessine (117), and oganesson (118)—confirmed in 2015 and named in 2016, completing the seventh period of the periodic table.^[7] Elements vary widely in abundance and stability: about 80% are metals, with rare gases like helium and neon being nonmetals, and unstable superheavy elements having half-lives as short as milliseconds.^[2] The periodic table thus not only catalogs these elements but also underpins advancements in chemistry, materials science, and nuclear physics.^[4]

Introduction

Definition and Characteristics

A chemical element is defined as a pure chemical substance composed of atoms with the same number of protons in the atomic nucleus.^[1] This definition emphasizes that elements are species of atoms sharing identical nuclear proton counts, forming the basis for their unique chemical identities. Each atom consists of a central nucleus containing protons and neutrons, surrounded by a cloud of electrons in discrete energy levels or shells.^[8] The atomic number, denoted as Z, represents the number of protons in the nucleus and serves as the defining feature that distinguishes one element from another.^[9] Isotopes are variants of an element that have the same atomic number but differ in the number of neutrons, thus varying in mass while retaining identical chemical properties.^[10] Additionally, some elements exhibit allotropes, which are different structural modifications of the same element in the same physical state, such as the diamond and graphite forms of carbon.^[11] In the atomic structure, protons determine the element's identity through Z, while the arrangement of electrons in outer shells governs its chemical reactivity and bonding behavior.^[12] Hydrogen exemplifies the simplest element, with Z=1, consisting of one proton in its nucleus and one electron, making it the lightest and most abundant element in the universe.^[9] In contrast, uranium, with Z=92, is the heaviest naturally occurring element available in significant quantities, featuring 92 protons and typically 146 neutrons in its most common isotope, uranium-238.^[13] These examples illustrate how atomic number scales from the minimal to the maximal in natural elements, influencing their stability and applications.

Role in Chemistry and Nature

Chemical elements serve as the fundamental building blocks of all matter, with 118 known elements forming the basis of every substance through various types of chemical bonding. Approximately 94 of these elements occur naturally on Earth, while the remaining are synthetic. These elements combine via ionic bonds, where electrons are transferred between metals and nonmetals; covalent bonds, involving shared electrons typically between nonmetals; and metallic bonds, which allow delocalized electrons to facilitate conductivity in metals.^[2][14][15] In biological systems, certain elements are indispensable for life processes. Carbon, hydrogen, oxygen, and nitrogen form the backbone of organic molecules essential to all known life forms, comprising the majority of biomass. Trace elements play critical roles as well; for instance, iron is a key component of hemoglobin, enabling oxygen transport in blood, while iodine is vital for thyroid hormone production, regulating metabolism.^[16][17][18] Elements are ubiquitous in the natural environment, shaping planetary composition and dynamics. Oxygen constitutes about 21% of Earth's atmosphere by volume, supporting respiration and combustion. Silicon, the second most abundant element in the crust at approximately 28% by mass, forms the structural basis of rocks and soils through silicate minerals. Elements also participate in geochemical cycles, such as the carbon cycle, which circulates carbon between the atmosphere, oceans, biosphere, and geosphere, influencing climate and sustaining life.^[19][20][21] Industrially, elements are engineered into alloys and advanced materials to enhance properties like strength and conductivity. Metals such as iron, chromium, and nickel are alloyed to create corrosion-resistant steels used in construction and machinery. In electronics, silicon and germanium serve as semiconductors, forming the core of transistors and integrated circuits that power modern computing devices.^[22][23]

Historical Context

Early Observations and Classifications

The concept of chemical elements originated in ancient philosophical traditions, where natural substances were categorized as fundamental building blocks of matter. Around 450 BCE, the Greek philosopher Empedocles proposed the four-element theory, positing that all matter consisted of earth, air, fire, and water, which combined and separated under the forces of love and strife.^[24] This framework was later adopted and expanded by Aristotle in the 4th century BCE, who integrated it into his natural philosophy by associating the elements with qualitative properties such as hot, cold, wet, and dry, influencing Western thought for centuries.^[25] Independently, in ancient Chinese philosophy, the wu xing system described five phases—wood, fire, earth, metal, and water—as dynamic, interdependent aspects of the natural world, used to explain cycles of change and interactions in cosmology and medicine.^[26] During the medieval and Renaissance periods, alchemical pursuits built on these ideas while emphasizing practical manipulations of substances. In the 16th century, the Swiss physician and alchemist Paracelsus shifted focus from transmutation to medicinal applications, classifying matter into the tria prima: salt (representing fixity), sulfur (combustibility), and mercury (volatility), which he viewed as the primary principles underlying all substances rather than the classical four elements.^[27] Alchemists of this era also isolated and worked with several metals known since antiquity, such as gold, silver, and copper, refining techniques for their purification and alloying, though these were not yet understood as distinct elements in the modern sense.^[28] The 18th century marked a transition toward empirical chemistry, with systematic observations challenging earlier theories. In 1766, Henry Cavendish isolated hydrogen by reacting metals with acids, describing it as "inflammable air" that produced water upon combustion, though he interpreted it through the lens of phlogiston theory.^[29] Joseph Priestley discovered oxygen in 1774 by heating mercuric oxide, noting its superior support for combustion compared to common air, and initially called it "dephlogisticated air."^[30] That same year, Carl Wilhelm Scheele produced chlorine by treating hydrochloric acid with manganese dioxide, mistaking it for a compound of oxygen but recognizing its distinct green-yellow gas and bleaching properties.^[31] These discoveries contributed to the debunking of the phlogiston theory, which posited a fire-like substance released during combustion; Antoine Lavoisier demonstrated through precise weighings that combustion involved oxygen gain rather than phlogiston loss, establishing conservation of mass.^[32] In 1789, Lavoisier published a foundational list of 33 substances he considered simple or elemental, including metals like iron and sulfur, non-metals such as oxygen and hydrogen, and even light and caloric (heat) as provisional elements, excluding compounds and laying groundwork for more rigorous classifications.^[33] These efforts highlighted the growing recognition of elements as irreducible components, paving the way for organized tabular systems in the following century.

Development of the Modern Periodic Table

In the early 19th century, Johann Wolfgang Döbereiner identified patterns among elements by grouping them into triads, where the atomic weight of the middle element was approximately the average of the other two, and the members exhibited similar chemical properties.^[34] For instance, lithium (Li), sodium (Na), and potassium (K) formed such a triad, with Na's atomic weight roughly midway between Li and K, and all three displaying alkaline reactivity.^[34] Döbereiner first noted these analogies in 1817 and formalized them in a 1829 publication, laying groundwork for recognizing regularities in elemental properties.^[34] By the mid-19th century, John Newlands advanced this idea with his law of octaves in 1865, arranging the 62 known elements in order of increasing atomic weight and observing that their properties repeated every eighth element, akin to musical octaves.^[35] Elements separated by intervals of seven in this sequence, such as hydrogen, fluorine, and chlorine, shared analogous characteristics like reactivity.^[35] Although limited to lighter elements and initially ridiculed, Newlands' work highlighted periodic repetition based on atomic weights.^[35] Dmitri Mendeleev independently formulated the periodic law in 1869, organizing elements by increasing atomic weight into a table where similar properties recurred at regular intervals, leaving gaps for undiscovered elements.^[36] He predicted properties for eka-aluminum (later gallium, atomic weight ~68-75, density ~5.9 g/cm³) and eka-silicon (later germanium, atomic weight ~72, density ~5.5 g/cm³), which were confirmed upon their discoveries in 1875 and 1886, respectively, validating his system.^[36] Around the same time, Julius Lothar Meyer independently developed a similar table in 1870, plotting atomic weights against volumes and grouping elements with analogous valencies across nine columns.^[37] In 1913, Henry Moseley refined the periodic table by demonstrating through X-ray spectroscopy that elemental order depends on atomic number (Z), the number of protons in the nucleus, rather than atomic weight, resolving discrepancies like the placement of argon and potassium.^[38] Moseley's law, ν = (3/4) R (Z - 1)² (1/1² - 1/2²) for Kα X-ray frequencies, confirmed Z as the fundamental ordering principle.^[38] That same year, Niels Bohr's atomic model incorporated quantized electron shells, explaining periodicity as arising from successive filling of these shells (K, L, M, etc.), with Z determining the number of electrons and thus chemical behavior.^[39] This integration solidified the modern framework, linking atomic structure to periodic trends.^[39]

Periodic Table Organization

Structure, Periods, and Groups

The periodic table organizes the chemical elements into a tabular arrangement that reveals patterns in their properties based on atomic structure. This layout consists of horizontal rows known as periods and vertical columns known as groups, reflecting the filling of electron shells and similarities in valence electron configurations, respectively.^[2][40] There are seven periods in the periodic table, corresponding to the principal quantum number of the outermost electron shell, which determines the energy level being filled. For instance, period 1 includes elements with electrons only in the 1s orbital, while period 2 involves the filling of the 2s and 2p orbitals. As atomic number increases across a period, electrons occupy orbitals within the same shell, leading to trends such as decreasing atomic radius and increasing electronegativity due to greater effective nuclear charge.^[41][42] Groups, or families, are the 18 vertical columns numbered 1 through 18 according to the IUPAC system, where elements share similar chemical behaviors arising from identical numbers of valence electrons in their outermost shells. Elements in the same group exhibit analogous reactivity patterns; for example, Group 1 (alkali metals) possess one valence electron, resulting in high reactivity as they readily lose it to form positive ions, while Group 17 (halogens) have seven valence electrons, making them strong oxidizing agents that gain one electron to achieve stability. These shared valence electron configurations underpin the predictive power of the table for chemical properties.^{[2][40][43][44]} The lanthanides and actinides form two 14-element series placed below the main body of the periodic table to accommodate the f-block elements, where the 4f and 5f orbitals are successively filled, respectively. These series span from cerium (atomic number 58) to lutetium (71) for the lanthanides and from thorium (90) to lawrencium (103) for the actinides, avoiding an excessively wide table. A key feature is the f-block contraction, particularly pronounced in the lanthanides, where atomic and ionic radii decrease more than expected across the series due to the poor shielding effect of 4f electrons, leading to increased effective nuclear charge and influencing properties like ion sizes in subsequent elements.^[45][46] This organization embodies the periodic law, which states that the physical and chemical properties of elements recur periodically when arranged in order of increasing atomic number, enabling the anticipation of undiscovered elements and their behaviors.^[47]

Blocks and Electron Configurations

The periodic table is divided into blocks based on the type of atomic orbital being filled by the valence electrons of the elements, reflecting the quantum mechanical description of electron arrangements in atoms. These blocks—s, p, d, and f—correspond to the subshells defined by the azimuthal quantum number $l$l, where $l = 0$l=0 for s, $l = 1$l=1 for p, $l = 2$l=2 for d, and $l = 3$l=3 for f subshells. The principal quantum number $n$n indicates the energy level or shell, with subshells filling in order of increasing $n$n and $l$l, influencing the periodicity of element properties such as reactivity and bonding. This organization arises from the wave-like nature of electrons, where each subshell can hold a maximum of $2(2l + 1)$2(2l+1) electrons, leading to capacities of 2 for s, 6 for p, 10 for d, and 14 for f subshells.^[48][49] The Aufbau principle governs the sequential filling of these orbitals in the ground state of atoms, stating that electrons occupy the lowest-energy available orbitals first, following the order 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, and 7p. This principle, combined with the Pauli exclusion principle (limiting each orbital to two electrons of opposite spin) and Hund's rule (maximizing unpaired electrons in degenerate orbitals), predicts most electron configurations accurately. However, exceptions occur due to enhanced stability from half-filled or fully filled subshells, such as in chromium (atomic number 24), which has the configuration [Ar] 4s¹ 3d⁵ instead of the expected [Ar] 4s² 3d⁴, prioritizing the half-filled 3d subshell for lower energy. Similarly, copper (atomic number 29) adopts [Ar] 4s¹ 3d¹⁰ over [Ar] 4s² 3d⁹ to achieve a full 3d subshell. These anomalies highlight how electron-electron repulsions and exchange energies can override strict Aufbau ordering in transition metals.^[50][51] The s-block encompasses groups 1 and 2 (alkali and alkaline earth metals), where valence electrons occupy the ns¹ or ns² configuration, making these elements highly electropositive and prone to losing one or two electrons to form cations. For instance, sodium (group 1) has [Ne] 3s¹, facilitating easy donation of its single valence electron in reactions. In contrast, the p-block (groups 13–18) features elements with np¹ to np⁶ valence configurations, leading to diverse behaviors from metallic (e.g., aluminum's [Ne] 3s² 3p¹) to nonmetallic (e.g., chlorine's [Ne] 3s² 3p⁵, which gains one electron to complete its octet). The d-block, comprising transition metals in groups 3–12, involves filling of (n-1)d¹⁻¹⁰ alongside ns¹⁻² valence electrons, resulting in variable oxidation states and colorful compounds due to partially filled d orbitals; iron, for example, is [Ar] 4s² 3d⁶. The f-block includes the lanthanides and actinides, with valence configurations of (n-2)f¹⁻¹⁴ and often ns², exhibiting complex magnetic properties and radioactivity in actinides due to poor shielding by f electrons.^[52][53][54] Valence electrons, those in the outermost shell or unfilled subshells, primarily dictate an element's chemical reactivity and bonding tendencies, as they participate in forming ions or sharing electrons. In s-block elements, the low ionization energies of ns electrons promote ionic bonding and high reactivity with water or oxygen, whereas p-block elements often form covalent bonds by sharing np electrons, enabling a wide range of molecular structures like those in organic chemistry. For d- and f-block elements, the involvement of multiple valence electrons from d or f subshells leads to coordination compounds and catalytic properties, underscoring how block-specific configurations underpin the periodic trends in electronegativity, metallic character, and compound formation.^[55][41]

Core List of Elements

Elements Ordered by Atomic Number

The chemical elements are ordered below by atomic number (Z), ranging from hydrogen (Z=1) to oganesson (Z=118), as recognized by the International Union of Pure and Applied Chemistry (IUPAC). This canonical sequence reflects the increasing number of protons in the atomic nucleus and serves as the foundation for the periodic table's arrangement. The table provides key properties for each element, including the chemical symbol, name, standard atomic weight (from IUPAC recommendations as of 2024, with values in brackets for synthetic elements indicating the mass number of the longest-lived isotope), density (at STP where applicable), melting point, boiling point (both in °C), phase at standard temperature and pressure (STP: 25°C and 1 atm), and Pauling electronegativity (dimensionless scale from 0.7 to 4.0, unavailable for noble gases and some synthetic elements). Physical properties like density, melting, and boiling points are compiled from standard reference data; for unstable synthetic elements (Z ≥ 93), many properties are estimated or not measured due to their short half-lives, often on the order of seconds to milliseconds—for instance, oganesson-294 has a half-life of approximately 0.7 ms. For synthetic elements (Z ≥ 93), many properties are estimated or unknown due to short half-lives.^{[56][57][58][59]}

Key Data Columns and Their Meanings

The atomic number, denoted as Z, represents the number of protons in the nucleus of an atom and serves as the fundamental identifier and ordering principle for chemical elements in the periodic table.^[61] This value determines the element's position and its unique chemical properties, as each increment in Z corresponds to a distinct element.^[2] The chemical symbol is a one- or two-letter abbreviation for the element's name, standardized by the International Union of Pure and Applied Chemistry (IUPAC), while the full name follows IUPAC-approved nomenclature.^[62] Symbols often derive from Latin or historical names, such as Fe for iron from the Latin ferrum, and for newly discovered elements, provisional systematic names like ununpentium (Uup) were used prior to official ratification in 2016.^[63] The atomic mass, or more precisely the standard atomic weight, is the weighted average of the masses of an element's naturally occurring isotopes, expressed in unified atomic mass units (u), where 1 u equals one-twelfth the mass of a carbon-12 atom.^[64] For example, carbon's value of 12.011 u reflects the abundance of carbon-13 alongside the dominant carbon-12 isotope.^[2] Synthetic elements, lacking stable isotopes, typically list estimated or mass number-based values with footnotes indicating variability due to short half-lives.^[65] Physical properties in element lists include density, measured in grams per cubic centimeter (g/cm³) at standard conditions; melting and boiling points, reported in kelvin (K) or degrees Celsius (°C) under 1 atm pressure; and phase of matter, classified as solid, liquid, or gas at 298 K (25°C) and 1 atm (101.325 kPa).^[57] These measurements adhere to International Union of Pure and Applied Chemistry (IUPAC) and National Institute of Standards and Technology (NIST) standards for consistency, with density often specified for the solid phase at 298 K unless otherwise noted.^[66] Electronegativity quantifies an atom's ability to attract electrons in a chemical bond and is commonly reported on the Pauling scale, a dimensionless measure ranging from approximately 0.7 (for cesium) to 4.0 (for fluorine).^[67] Introduced by Linus Pauling, this scale derives from bond energy differences and helps predict bond polarity and reactivity, with values influencing electron configurations that underpin periodic trends.^[67]

Synthetic Elements

Methods of Synthesis

The synthesis of elements beyond uranium (atomic number 92), known as transuranic or superheavy elements, primarily relies on nuclear reactions facilitated by particle accelerators. These methods involve bombarding target nuclei with high-energy particles to induce fusion or capture processes, producing new isotopes that decay rapidly. Early transuranics were created through neutron capture or charged-particle bombardment, while superheavy elements (atomic numbers 104 and above) are typically synthesized via heavy-ion fusion reactions followed by neutron evaporation.^[68][69] Particle accelerators, such as cyclotrons and linear accelerators, are essential for achieving the required energies, often in the range of several MeV per nucleon, to overcome the Coulomb barrier between nuclei. Cyclotrons, which accelerate particles in a spiral path using a magnetic field, were pivotal in the initial discoveries of transuranic elements at the University of California, Berkeley. For instance, in December 1940, Glenn T. Seaborg and colleagues used a 60-inch cyclotron to bombard uranium-238 with deuterons, producing neptunium-238, which beta-decayed to yield the first atoms of plutonium (element 94). This marked the dawn of synthetic transuranics, with subsequent elements like americium and curium following similar charged-particle or neutron-induced reactions in cyclotrons and early reactors. Linear accelerators, which propel ions in a straight line via radiofrequency fields, have since become standard for superheavy element production due to their ability to handle heavier beams at higher intensities.^[70][68] Fusion reactions dominate the creation of superheavy elements, where accelerated heavy ions collide with heavy target nuclei to form a compound nucleus that evaporates neutrons to stabilize. A key approach, known as "cold fusion," uses projectiles like lead or calcium on bismuth or deformed targets to minimize excitation energy and maximize survival probability. At the GSI Helmholtz Centre in Darmstadt, Germany, elements 107 (bohrium) through 112 (copernicium) were synthesized from the 1980s to the 1990s using the velocity filter SHIP (Separator for Heavy Ion Reaction Products) with lead-208 or calcium-48 beams on bismuth-209 or other heavy targets, for example, ^{209}Bi(^{64}Ni,n)^{272}Rg for roentgenium (element 111). These experiments confirmed single-neutron evaporation channels, with cross-sections on the order of picobarns, requiring weeks of irradiation to detect a few atoms.^[71][72] Notable milestones highlight international efforts in superheavy synthesis. In 2004, researchers at Japan's RIKEN Nishina Center used the RIKEN Linear Accelerator (RILAC) to fuse zinc-70 with bismuth-209, producing nihonium (element 113) as ^{278}113 via a 1n-evaporation channel, with the first decay chain observed on July 23. Similarly, in 2006, the Joint Institute for Nuclear Research (JINR) in Dubna, Russia, employed a calcium-48 beam on californium-249 at their U-400 cyclotron to synthesize oganesson (element 118) through the reaction ^{249}Cf(^{48}Ca,3n)^{294}Og, detecting three decay chains that confirmed its production. These "hot fusion" reactions, using actinide targets, enable access to neutron-richer isotopes compared to cold fusion.^[73][74] Verification of these discoveries involves rigorous international collaboration through the International Union of Pure and Applied Chemistry (IUPAC) and the International Union of Pure and Applied Physics (IUPAP). A joint working group reviews experimental data, including decay chains and cross-sections, often over several months to years; for example, the confirmation of elements 113–118 in 2015–2016 followed independent reproductions and an approximately eight-month peer review process to ensure reproducibility and rule out contaminants. This process, established since the 1990s for transfermium elements, ensures only validated claims lead to naming rights.^[75][76] Recent advances include a 2025 experiment at Oak Ridge National Laboratory producing livermorium via ^{50}Ti beams on ^{244}Pu targets, enhancing prospects for synthesizing element 119. As of November 2025, no elements beyond oganesson (Z=118) have been officially recognized by IUPAC.^[77]

Properties and Stability

Synthetic elements, particularly those beyond uranium (transuranic elements with atomic numbers Z ≥ 93), exhibit extreme instability due to their high nuclear charge and neutron-proton imbalance, leading to very short half-lives that span from microseconds to centuries. For instance, hassium (Z=108), ^{270}Hs has a half-life of about 3.6 seconds, decaying primarily via alpha emission to seaborgium-266.^[78] In contrast, longer-lived isotopes like americium-241 (Z=95) possess a half-life of 432.2 years, though even this is insufficient for natural persistence without continuous synthesis. Alpha decay dominates as the primary decay mode in these transuranic elements, where the nucleus emits an alpha particle (helium-4 nucleus) to reduce its mass and achieve greater stability, often followed by subsequent decays in a chain until reaching a more stable lighter nucleus. Relativistic effects become increasingly pronounced in superheavy elements (Z > 100), where the strong electrostatic attraction from the nucleus accelerates inner electrons to speeds approaching the speed of light (c), causing a relativistic mass increase that contracts s and p_{1/2} orbitals while expanding p_{3/2}, d, and f orbitals. This distortion alters chemical bonding and reactivity; for example, oganesson (Z=118), expected to behave as a noble gas homologue of radon, is predicted to be more reactive due to these effects destabilizing its closed-shell configuration and enhancing spin-orbit coupling, potentially allowing compound formation under specific conditions. Such modifications challenge traditional periodic trends, making superheavy chemistry deviate from lighter analogues. Theoretical models predict an "island of stability" for superheavy nuclei around Z = 114–126 and neutron numbers N ≈ 184, where enhanced shell closures could extend half-lives to seconds or even minutes, contrasting with the rapid decay of currently synthesized isotopes. Elements in this region, such as flerovium (Z=114), are anticipated to share some properties with group 14 homologues like lead, including metallic character, but with heightened volatility—flerovium is expected to be the most volatile metal in the group, depositing on gold surfaces at room temperature while forming bonds stronger than those of noble gases under forcing conditions. These predictions rely on relativistic quantum calculations, as direct observation remains elusive. Experimental characterization of synthetic elements is severely limited by their fleeting existence, with detections typically occurring at the single-atom level through analysis of alpha decay chains using gas-phase chromatography or recoil separators. No bulk samples have been produced, precluding traditional spectroscopic or thermodynamic measurements; instead, properties are inferred from the rapid succession of decay events and adsorption behaviors on detector surfaces.

Occurrence and Applications

Natural Abundance and Sources

The cosmic abundance of chemical elements is dominated by hydrogen and helium, which together constitute approximately 97.3% of the ordinary matter in the universe by mass, with hydrogen accounting for about 73.5% and helium for 23.8%; oxygen follows as the third most abundant at roughly 1%. These distributions arise primarily from Big Bang nucleosynthesis, which produced the primordial hydrogen and helium in the early universe, while heavier elements like oxygen are synthesized through stellar nucleosynthesis processes such as fusion in the cores of stars.^[79][80] On Earth, the crustal composition differs markedly from cosmic abundances due to planetary differentiation and geological processes. Oxygen is the most prevalent element in the Earth's crust at 46.6% by weight, followed by silicon at 27.7% and aluminum at 8.1%, forming the basis of most silicate minerals. The mantle is enriched in magnesium and iron, while the core consists primarily of iron (about 85-90%) with significant nickel (5-10%), reflecting the segregation of denser metals during Earth's formation.^[81] In the atmosphere, nitrogen comprises 78.08% by volume and oxygen 20.95%, with the remainder consisting of trace gases like argon and carbon dioxide; these proportions result from volcanic outgassing and biological cycling over billions of years. The hydrosphere, particularly seawater, features sodium and chloride as the dominant ions, with chloride at about 18,980 mg/L and sodium at 10,556 mg/L, constituting over 85% of the total dissolved salts by weight.^[19][82] Biological systems concentrate a subset of elements essential for life, with hydrogen, oxygen, carbon, and nitrogen making up approximately 96% of the biomass by mass in living organisms, primarily through incorporation into water, proteins, carbohydrates, and nucleic acids. Rare earth elements, such as cerium and lanthanum, occur in trace amounts within biological contexts, including recent discoveries of nanoscale monazite formation in living ferns as of November 2025, but are more prominently associated with minerals like monazite, which can contain up to 50-60% rare earth oxides, including these two as major components.^[83][84][85] Certain heavy elements on Earth are primordial, incorporated during planetary accretion, or radiogenic, produced by decay chains. Uranium and thorium, for instance, originated from rapid neutron-capture processes in supernovae preceding the solar system's formation, with their current abundances sustained by long half-lives (uranium-238 at 4.5 billion years and thorium-232 at 14 billion years). Extinct radionuclides like plutonium-244, with a half-life of 81.3 million years, were present in the early solar system but have since decayed completely, leaving isotopic signatures in meteorites as evidence of their prior existence.^[86][87]

Industrial Uses and Economic Importance

Iron is a cornerstone of the global metals industry, primarily through its role in steel production, which reached 1.89 billion metric tons in 2024 and supports infrastructure, automotive, and construction sectors.^[88] Steel's versatility drives economic activity, with iron ore mining and processing contributing significantly to commodity trade. Aluminum, refined from bauxite via the Bayer process, is essential in aerospace for its lightweight strength, enabling fuel-efficient aircraft designs; the global aluminum market is projected to reach $355 billion by 2025.^[89] Rare earth elements, particularly neodymium used in high-performance permanent magnets for electric motors and wind turbines, face supply chain vulnerabilities, with China controlling approximately 70% of global mining output and 90% of refining capacity as of 2025.^[90] The rare earth metals market was valued at $6.4 billion in 2024, underscoring their critical economic role in renewable energy technologies.^[91] Among non-metals, carbon's allotropes find diverse industrial applications; graphene enhances electronics and composites for its superior conductivity and strength, with the global graphene market expected to grow from $941.1 million in 2025 to $8.3 billion by 2032.^[92] Synthetic diamonds serve in cutting tools, abrasives, and semiconductors due to their hardness and thermal properties, contributing to a diamond industry valued at around $110 billion in 2025.^[93] Silicon powers the electronics sector through semiconductors and photovoltaic (PV) solar cells; global polysilicon production capacity for solar reached 3.35 million tons by early 2025, supporting over 1,500 GW of potential PV module output.^[94] Oxygen is vital for industrial processes like steelmaking (enhancing combustion efficiency) and oxy-fuel welding, as well as medical applications such as respiratory therapy; its supply chain involves cryogenic air separation, with moderate economic risks due to widespread production.^[95] In the energy and radioactive elements category, uranium serves as nuclear fuel, with global production estimated at 55,000 to 65,000 tons annually to meet reactor demands projected at 68,920 tons in 2025.^[96][97] Lithium has surged in importance for lithium-ion batteries amid the electric vehicle (EV) boom since 2020, with EV sales exceeding 17 million units in 2024 and driving battery demand; the lithium-ion battery market is forecasted to expand from $54.4 billion in 2023 to $182.5 billion by 2030.^[98][99] Economic aspects of these elements highlight supply chain dependencies and sustainability efforts. China's dominance in rare earth processing (over 90%) exemplifies geopolitical risks, prompting diversification initiatives.^[100] Copper recycling achieves a global end-of-life rate of about 40%, aiding supply stability for wiring and renewables.^[101] Gold, used in electronics and as a financial asset, traded at approximately $4,000 per ounce in November 2025, reflecting its role in investment portfolios amid economic uncertainty.^[60]