Periodic table

• 1. Introduction

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  The periodic table arranges the chemical elements into rows called periods and columns called groups. Hidden inside that grid is a law. Arrange the elements in order of their atomic numbers, and an approximate recurrence of their properties appears. That single idea, the periodic law, turned a list of substances into a map.

  In 1869 a Russian chemist named Dmitri Mendeleev built the first version of this map that the world accepted. Not every element was known yet, so his table had gaps. Mendeleev did something audacious with those empty spaces. He used the periodic law to predict properties of elements no one had ever seen.

  How can a chart of rows and columns predict the unknown? Why does metallic character grow as you move down a column and toward the left? Why do some elements refuse to sit comfortably in any group? And why do scientists believe that out past the edges of the known table, the patterns themselves may break down? The answers run from the inside of the atom to laboratories that build elements that nature never made.

• 2. What The Grid Counts

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  Each chemical element carries a unique atomic number, written Z, from the German Zahl, meaning number. That number is simply how many protons sit in the nucleus. Hydrogen is atomic number 1, helium is 2, lithium is 3, and so on up the line. Each name shortens to a one- or two-letter symbol: H, He, Li.

  Neutrons complicate the picture without changing an element's identity. Atoms with the same proton count but different neutron counts are isotopes of the same element. Carbon shows this plainly. All its atoms have six protons, most have six neutrons, about one per cent carry seven, and a very small fraction carry eight. The table never splits isotopes apart. They always sit together under a single element.

  A new period begins when a new electron shell takes its first electron. The columns come from electron configuration. Oxygen, sulfur, and selenium share a column because each has four electrons in its outermost p-subshell. Elements close together on the table tend to behave alike, which means the properties of one element can often be guessed from its neighbours.

  Of the 118 known elements, the first 94 occur naturally on Earth. The remaining 24, americium through oganesson, exist only when synthesized in laboratories. Even among the natural 94, a few were built in the lab before anyone realized they also existed in nature: technetium, promethium, astatine, neptunium, and plutonium.

• 3. Survivors And Ghosts

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  Bismuth holds an almost-stable isotope with a half-life of 2.01 by 10 to the 19th power years, more than a billion times the age of the universe. That is the kind of timescale that separates the elements that have lasted from the ones that vanish. Of the 94 natural elements, eighty have a stable isotope, and bismuth adds one more on the strength of that near-eternal half-life.

  Thorium and uranium tell a different survival story. Their decaying isotopes have half-lives comparable to the age of the Earth itself. Together with the stable elements and bismuth, they make up the 83 primordial elements that have endured since the Earth formed.

  The remaining eleven natural elements are ghosts that keep reappearing. They decay too quickly to survive on their own. Their trace presence rests on constant regeneration as intermediate products in the decay chains of thorium and uranium. Some elements resist being seen at all. Francium has only ever been photographed as light emitted from microscopic quantities. No element heavier than einsteinium has been observed in macroscopic amounts in pure form, and the same is true of astatine. Every one of the 24 artificial elements is radioactive.

• 4. Shells And Orbitals

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  An electron can be pictured as inhabiting an atomic orbital, a region describing where it is likely to be found. Its energy is quantised, taking only discrete values. Electrons obey the Pauli exclusion principle, meaning no two can occupy the same state. This sorts them into shells, then subshells, then orbitals, each orbital holding up to two electrons distinguished by a spin labelled up or down.

  The first shell holds a single spherical s orbital and just two electrons. The second shell adds three dumbbell-shaped p orbitals, raising its capacity to eight. The third shell, with one s, three p, and five d orbitals, holds eighteen. The fourth shell adds seven f orbitals and reaches thirty-two. The pattern follows 2 times n squared.

  Only the outermost electrons, called valence electrons, carry enough energy to break free and join chemical reactions. The rest are core electrons, bound too tightly to participate. This split between core and valence is what makes the table's chemistry repeat. The order in which subshells fill follows the Aufbau principle, also known as the Madelung or Klechkovsky rule, after Erwin Madelung and Vsevolod Klechkovsky. Electrons enter orbitals in order of increasing n plus the azimuthal quantum number, and when two orbitals tie, the one with lower n fills first.

• 5. Building The Rows

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  Hydrogen places its single electron in the lowest orbital, written 1s to the 1. Helium adds a second, filling the first shell as 1s squared. From lithium onward the full 1s subshell becomes a core shell, so lithium's third electron starts the second shell at 2s. The march continues through beryllium, then boron, carbon, and nitrogen filling the three 2p orbitals one at a time, following Hund's rule that orbitals fill singly before doubling up.

  Sodium opens the third shell, and its configuration abbreviates to neon's plus 3s to the 1. The outer structures of sodium through argon mirror those of lithium through neon. That echo, repeating at regular intervals of atomic number, is the periodicity the whole table illustrates.

  Scandium introduces a wrinkle. The 4s and 3d subshells sit at nearly equal energy and compete, so the 3d orbitals do not fill in a clean sequence. At chromium the atom prefers a 3d to the 5, 4s to the 1 arrangement, and copper takes 3d to the 10, 4s to the 1. Both break the Madelung rule. These anomalies carry no chemical significance, since the configurations sit so close in energy that a nearby atom can shift the balance, so the table simply ignores them. The d-block elements, filling an inner shell, are the transition metals, and from gallium onward the filled 3d orbitals retreat into the core.

• 6. The Helium Problem

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  Helium is the only element that routinely sits in a position that contradicts its own electronic structure. It has two outer electrons, while the other noble gases have eight, and it is an s-block element while they are all p-block. Yet its full outer shell and its inertness match group 18, so that is almost always where it goes. A 1988 proposal to move helium to group 2 was rejected by IUPAC.

  The case for group 2 has not died. Solid helium crystallises in a hexagonal close-packed structure, matching beryllium and magnesium rather than the other noble gases. Theoretical work suggests helium may be slightly less inert than neon and could form a compound written HeO bonded with two LiF units, echoing an analogous beryllium compound with no neon counterpart. Such molecules would likely only survive near 10 K.

  Hydrogen creates its own dispute. Like the alkali metals it has one outer electron and usually loses it, and it can displace some metals from their salts. But it forms a diatomic nonmetallic gas, unlike the reactive solid alkali metals, and it forms hydrides by gaining an electron, which links it to the halogens. Hydrogen is neither strongly oxidizing nor strongly reducing and does not react with water. The electronic placement in group 1 dominates, though some tables float it apart from every group. The chemist and philosopher Eric Scerri criticized that floating option, arguing it implies hydrogen stands above the periodic law itself.

• 7. The Group 3 Dispute

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  Lev Landau and Evgeny Lifshitz, in 1948, judged it incorrect to group lutetium as an f-block element, because the 4f shell completes its filling at ytterbium. They stopped short of also removing lanthanum from the d-block. The puzzle of where the f-block truly begins sat unresolved.

  Jun Kondo took the next step in 1963, realizing that lanthanum's low-temperature superconductivity revealed the activity of its 4f shell. In 1965 David C. Hamilton tied that observation to the table's structure, arguing the f-block should run from lanthanum to ytterbium and from actinium to nobelium. William B. Jensen brought the matter to wide attention in 1982, and IUPAC reports from 1988 and 2021 supported placing lutetium and lawrencium in group 3. The older arrangement persists mainly because many textbook writers are unaware of the issue.

  The deciding test is which orbitals can do chemistry. Lanthanum and actinium, like thorium, have valence f orbitals that can fill in chemical environments, while lutetium and lawrencium keep their f-shells locked in the core. That makes the bond between yttrium and lutetium a primary relationship, sharing both valence electron count and orbital type, rather than the looser link between yttrium and lanthanum. A third form leaves the spaces below yttrium empty, but that makes the f-block fifteen elements wide even though only fourteen electrons fit in an f-subshell.

• 8. Trends And The Relativistic Edge

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  Arthur Haas published the first calculated estimate of hydrogen's atomic radius in 1910, landing within a factor of ten of the accepted Bohr radius of about 0.529 angstroms. He used a single-electron model built on J. J. Thomson's 1904 plum-pudding atom. Atomic radii generally shrink moving left to right across a period as nuclear charge rises, and grow moving down a group as outer electrons occupy higher shells.

  The first row of each block runs abnormally small, an effect called kainosymmetry, because the 1s, 2p, 3d, and 4f subshells have no inner analogues to repel them. This pushes the small 2p elements toward multiple bonding that the larger 3p elements avoid. It also drives an even-odd zigzag known as secondary periodicity. Phosphorus and antimony readily reach the plus 5 oxidation state, while nitrogen, arsenic, and bismuth prefer plus 3.

  Ionisation energy, the cost of removing an electron, rises left to right and bottom to top, peaking at each period's noble gas. Electron affinity, the energy released on gaining an electron, peaks at the halogens, since the noble gases have no room for more. The longest-lived He minus level autodetaches in about 359 microseconds. Electronegativity follows the same shape on the Pauling scale, from fluorine at 4.0 down to caesium at 0.79.

  When nuclei grow highly charged, special relativity enters. Relativistic effects explain why gold is golden and why mercury is liquid at room temperature, and spin-orbit interaction splits the p subshell so that thallium and lead shrink while bismuth through radon expand. These effects are expected to grow very strong in the late seventh period, potentially collapsing periodicity. Electron configurations are clearly known only up to element 108, hassium, and experimental chemistry beyond it has reached only elements 112 through 115. The seventh row was completed when tennessine was synthesized in 2009, and its last elements were named in 2016.