Chemical element

• 1. Introduction

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  A chemical element is a species of atom defined by one number: how many protons sit in its nucleus. Oxygen has 8 protons, so its atomic number is 8. Carbon has 6, so its atomic number is 6. Change that proton count, and you have changed the element itself. This single rule governs almost all the matter in the universe. Yet the idea took millennia to arrive. Ancient thinkers spoke of earth, water, air, and fire. By November 2016, the International Union of Pure and Applied Chemistry recognized 118 elements. How did a vague notion of indivisible stuff become a precise count of proton numbers? Who built the table that now hangs in every classroom, and why do most of those elements never appear on Earth at all? The answers run from a Greek dialogue written around 360 BCE to laboratories that conjure atoms lasting fractions of a second.

• 2. The Proton Count

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  All carbon atoms contain 6 protons in their nucleus, and that fact alone makes them carbon. The number of protons sets the electric charge of the nucleus. That charge fixes how many electrons surround a neutral atom, and those electrons occupy orbitals that decide the atom's chemistry. Neutrons can vary without changing the identity. Carbon atoms may carry 6, 7, or 8 neutrons, giving carbon-12, carbon-13, and carbon-14. Natural carbon is about 98.9% carbon-12, about 1.1% carbon-13, and roughly one atom per trillion of carbon-14. These isotopes behave almost identically in chemical reactions because each still has six electrons. The neutron count barely touches chemical behavior, with hydrogen as a notable exception where the kinetic isotope effect is significant. That is why atomic number, not mass, identifies an element. Neutrons still matter for a different reason. The neutron-proton ratio determines whether a nucleus holds together against the mutual repulsion of its protons. Light elements up through oxygen can stay stable with equal neutrons and protons, but lead requires about 3 neutrons for every 2 protons. The strong nuclear force binds the bundle together over short range, and the right balance of neutrons keeps the whole thing from flying apart.

• 3. Born In Stars

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  Almost all baryonic matter in the universe is made of elements, with rare exceptions such as neutron stars. The lightest two, hydrogen and helium, came from Big Bang nucleosynthesis in the first 20 minutes of the universe, in a ratio of about 3:1 by mass, or 12:1 by number of atoms. Tiny traces of lithium and beryllium followed. No element heavier than boron formed in the Big Bang, and the primordial mix settled at roughly 75% hydrogen, 25% helium, and 0.01% deuterium. Nuclear fusion inside stars built the next wave. Stellar nucleosynthesis produces every element from carbon up to iron. The Sun's energy comes from fusing hydrogen into helium, slowly enriching its core with helium. In stars far more massive than the Sun, the alpha process forges heavier nuclei until iron is reached. There the buildup stalls, because the binding energy of a nucleus peaks at iron and nickel. Past that point fusion absorbs energy rather than releasing it, and an inert iron core forms. Elements heavier than iron, including uranium and plutonium, are made by explosive nucleosynthesis in supernovae and neutron star mergers. The light trio lithium, beryllium, and boron come mostly from cosmic ray spallation, the fragmentation of carbon, nitrogen, and oxygen struck by cosmic rays. Of the 118 known elements, the first 94 occur naturally, detected as primordial nuclides or as products of uranium and thorium. The Milky Way's most abundant element is hydrogen, at 739,000 parts per million by mass, with helium next at 240,000.

• 4. The Edge Of Stability

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  Every element has radioactive isotopes, though many never appear in nature because their half-lives are too short. Radioisotopes decay by alpha decay, beta decay, inverse beta decay, and for the heaviest, spontaneous fission. Isotopes with even numbers of protons or neutrons tend to be more stable, because like particles pair up with opposite spins and raise the binding energy. Most naturally occurring elements, 54 of the 94, have more than one stable isotope. Only 26 are monoisotopic, with exactly one stable isotope. Tin, element 50, holds the record at 10 stable isotopes. Across the 80 stable elements, the mean is 3.1 stable isotopes each. Technetium, element 43, and promethium, element 61, have no stable isotopes at all, even though they sit below element 82. Elements 83 through 94 are unstable enough that decay can be detected in every isotope. Yet bismuth at 83, thorium at 90, and uranium at 92 carry isotopes long-lived enough to survive from before the Solar System formed. Bismuth-209 has the longest known alpha decay half-life of any nuclide, far longer than the age of the universe, and is treated almost as a stable element. The 24 heaviest elements, those beyond plutonium, decay so fast that they are not found on Earth and must be synthesized. Five of them, from americium to einsteinium, have been spotted in the spectrum of Przybylski's star. Technetium was the first element synthesized that was thought not to occur naturally, made in 1937, though traces were later found in nature.

• 5. Electron Shells

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  Electrons confined to an atom may take only certain discrete energy levels, a restriction called quantization. Each electron's wave is held like a standing wave with a specific wavelength, and each whole number of wavelengths yields orbitals that map the electron's charge. Every orbital holds a pair of electrons. These orbitals gather into shells, each labeled by a principal quantum number marking its energy. The shells fill in a fixed order. Ordinary hydrogen, the simplest atom, has one proton and one electron, which sits in the first shell, called K, in subshell 1s. Helium adds a second electron to that same orbital, pairing their spins and completing the shell. Lithium's third electron cannot fit, so it moves to the second shell, called L, in subshell 2s. The shells can hold 2, 8, 18, 32, and so on, a sequence that draws the rows of the periodic table. The electron configuration of an atom governs how it bonds with neighbors. That is why the structure of the periodic table follows directly from the way electrons stack into shells, with each row sharing the same number of shells and each column dominated by the outermost electron's orbital.

• 6. Many Forms, Many States

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  Only bromine and mercury are liquid at 0 degrees Celsius and 1 atmosphere of pressure. Caesium and gallium are solid there but melt at 28.4 C and 29.8 C. Most elements are solid at standard temperature and pressure, and several are gases. Gallium spans the widest range between melting and boiling, reaching a boiling point of 2204 C. Elements that boil above 2000 are called refractory, while those that vaporize easily are volatiles. Helium is stranger still. It stays liquid even at absolute zero under atmospheric pressure, so it has a boiling point but no melting point in conventional tables. Beyond their states, many elements take multiple structural forms called allotropes. Carbon, sulfur, phosphorus, oxygen, and nitrogen are known for this. Carbon alone appears as diamond, with a tetrahedral structure; graphite, with stacked hexagonal layers; graphene, a single very strong layer; nearly spherical fullerenes; and carbon nanotubes. The familiar carbon allotropes carry different densities: amorphous carbon at 1.8 to 2.1, graphite at 2.267, and diamond at 3.515 grams per cubic centimetre. Solid elements settle into crystalline forms drawn from seven crystal families, including cubic, hexagonal, and orthorhombic. Over 30 elements crystallize in the cubic form, and 40% form close-packed crystals. Under a planetary interior's pressure, new crystalline forms appear; seven dense classes of silicon crystals can emerge at pressures from 1 to 100 at room temperature.

• 7. Naming The Atoms

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  The current system of chemical notation was invented by Jons Jacob Berzelius in 1814. He drew his symbols from the Latin names of elements, abbreviations meant to work across all languages and alphabets. In most cases the Latin root matches the modern English name, but eleven do not. Iron is Fe from ferrum, mercury Hg from hydrargyrum, gold Au from aurum, silver Ag from argentum, lead Pb from plumbum, tin Sn from stannum, copper Cu from cuprum, and antimony Sb from stibium. Sodium is Na from natrium, potassium K from kalium, and tungsten W from wolframium. Berzelius first suggested So and Po for sodium and potassium, then changed them to Na and K the same year. Naming new elements once belonged to their discoverers, and the choices grew national. Lutetium was named after Paris, and German chemists, reluctant to yield to the French, often called it cassiopeium instead. The British discoverer of niobium first named it columbium for the New World, a usage common in American publications until international standardisation in 1950. That friction pushed the rules toward a single authority. In 1947 an IUPAC conference decided that the names and symbols of new elements would be settled by IUPAC. Discoverers may still suggest a name, but the official choice, ancient or modern, rests with the organization. IUPAC favors British spellings such as aluminium and caesium, alongside the American sulfur, and treats element names as common nouns, so californium and einsteinium stay lowercase even though their symbols Cf and Es are capitalized.

• 8. From Fire To Proton

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  The word elements, stoicheia, was first used by Plato around 360 BCE in his dialogue Timaeus. He held that the elements Empedocles had introduced a century earlier were small polyhedral forms: a tetrahedron for fire, octahedron for air, icosahedron for water, and cube for earth. Around 350 BCE Aristotle kept the term, added a fifth element, aether, for the heavens, and defined an element as a body into which others decompose but which cannot itself be divided further. The chemical era began to pull the idea apart. In 1661, in The Sceptical Chymist, Robert Boyle argued against fixing a predetermined number of elements and favored matter built from irreducible corpuscles he also called atomes. He showed that gold reacts with aqua regia and mercury with various acids, then can be recovered, just as elements should be. In 1724 the minister and logician Isaac Watts listed the elements then recognised, mixing classical and chemical ideas, and noted that chemists did not all agree. The modern count grew steadily. Antoine Lavoisier's 1789 Elements of Chemistry held 33 elements, including light and caloric. By 1818 Berzelius had determined atomic weights for 45 of the 49 then-accepted elements. Dmitri Mendeleev placed 63 elements in his 1869 periodic table, built to show recurring trends. The decisive turn came in 1913, when Henry Moseley found that nuclear charge is the physical basis of the atomic number. The proton count replaced atomic weight as the defining trait. Francium, the last naturally occurring element to be found, was discovered in 1939 by Marguerite Perey. In 1955 element 101 was named mendelevium, honoring the man who first arranged the elements in order.