The standard periodic table arranges all 118 known chemical elements sequentially by their atomic number—the total number of protons found in the atomic nucleus—reading from left to right and top to bottom.¹² New horizontal rows are started, and blank spaces are inserted across the upper rows, so that elements with analogous chemical behaviors align vertically.

The horizontal rows are designated as periods and numbered from 1 to 7.¹³ A new period begins each time an electron initiates the filling of a new principal energy shell.¹⁴ Period 1 contains only two elements (hydrogen and helium), while periods 2 and 3 contain eight elements each. Periods 4 and 5 each comprise 18 elements, and periods 6 and 7 each contain 32 elements.¹⁵ To keep the printed width of the chart manageable for standard pages and wall posters, the 18-column medium-long form extracts 14 elements each from periods 6 and 7—the lanthanoids (or lanthanides) and actinoids (or actinides)—and places them in two detached rows beneath the main body of the table.¹⁶ When these elements are incorporated directly into the main grid between the s-block and d-block, the table expands into its continuous 32-column long form.¹⁷

The vertical columns are called groups, numbered from 1 to 18 from left to right under the naming convention established by the International Union of Pure and Applied Chemistry (IUPAC). Elements within a given group generally possess the same number of valence electrons in their outermost shells, giving rise to distinct chemical families. For example, Group 1 (excluding hydrogen) consists of the highly reactive alkali metals, Group 2 holds the alkaline earth metals, Group 17 contains the halogens, and Group 18 comprises the largely unreactive noble gases.¹⁸

The table is also partitioned into four distinct blocks—named the s-block, p-block, d-block, and f-block—based on the subshell in which the "last" electron resides according to the building-up (Aufbau) principle.¹⁹ The s-block encompasses Groups 1 and 2 plus helium; the p-block spans Groups 13 through 18 (except helium); the d-block corresponds to the transition metals in Groups 3 through 12; and the f-block comprises the detached inner transition series at the bottom of the standard chart.²⁰

While nineteenth-century chemists organized the periodic table empirically through observed chemical valences and atomic masses, the underlying mechanism governing periodicity is rooted in quantum mechanics and the electronic configuration of atoms. According to the Pauli exclusion principle, no two electrons in a single atom can share the exact same set of four quantum numbers.²¹ Consequently, electrons occupy discrete atomic orbitals arranged in shells and subshells of increasing energy.

The order in which these orbitals are filled is described approximately by the Madelung rule (often referred to as the $n + \ell$ rule, part of the Aufbau principle), where $n$ represents the principal quantum number (shell) and $\ell$ represents the azimuthal quantum number (subshell: $s=0, p=1, d=2, f=3$). Orbitals with a lower $n + \ell$ sum fill first; when two orbitals share the same $n + \ell$ value, the one with the lower $n$ value fills first. Because the $s$, $p$, $d$, and $f$ subshells contain one, three, five, and seven orbitals respectively—and each orbital holds at most two electrons of opposite spin—the subshells accommodate maximums of 2, 6, 10, and 14 electrons.

This sequential capacity directly dictates the lengths of the periods in the periodic table: 2 elements in the first period ($1s$), 8 in the second and third ($2s2p$ and $3s3p$), 18 in the fourth and fifth (adding the $3d$ and $4d$ subshells), and 32 in the sixth and seventh (adding the $4f$ and $5f$ subshells). When an outer $p$ subshell is completely filled with six electrons (or two $1s$ electrons in the case of helium), the atom achieves a highly stable closed-shell configuration characteristic of the Group 18 noble gases. The very next element must place an electron into a new, higher principal shell ($s$ orbital), resetting the periodic cycle with a reactive Group 1 alkali metal.

Because the arrangement of the table mirrors progressive changes in nuclear charge and electron shell structure, physical and chemical properties vary systematically across periods and down groups. These recurring patterns are known as periodic trends.

Atomic radius—a measure of the size of an atom—generally decreases from left to right across a period and increases from top to bottom down a group. Moving across a period, protons are added to the nucleus while electrons are added to the same principal energy shell. The increasing positive nuclear charge pulls the electron cloud closer to the nucleus without a commensurate increase in electron-electron shielding, contracting the atom. Moving down a group, each successive element adds an entirely new principal electron shell, placing the valence electrons further from the nucleus and increasing the atomic radius despite the higher nuclear charge.

Ionization energy (the energy required to remove an electron from a gaseous atom) and electronegativity (the tendency of an atom to attract shared electrons in a chemical bond) exhibit the inverse pattern. Both generally increase from left to right across a period as the shrinking atomic radius and rising effective nuclear charge bind valence electrons more tightly. Conversely, both generally decrease moving down a group because the outermost electrons are situated further from the nucleus and are increasingly shielded by inner core electron shells. Electron affinity—the energy change when a neutral atom gains an electron—similarly tends to become more negative (indicating a stronger attraction for an added electron) toward the upper right of the table, peaking at the halogens before dropping sharply at the closed-shell noble gases.

These electronic parameters govern the macroscopic transition between metallic and nonmetallic character. Metals, which dominate the left and center of the periodic table (the s-, d-, and f-blocks, plus the lower-left p-block), are characterized by low ionization energies and low electronegativities; they readily lose valence electrons to form positive cations and metallic bonds. Nonmetals, clustered in the upper right of the p-block, possess high ionization energies and electronegativities, tending to gain or share electrons in covalent bonds. A diagonal band of elements between these regions—including boron, silicon, germanium, arsenic, antimony, and tellurium—exhibits intermediate properties and is traditionally classified as metalloids or semimetals.³ Consequently, metallic character increases moving down a group and from right to left across a period, while nonmetallic character increases moving from the bottom left to the top right.

Despite the ubiquity of the 18-column periodic table, chemists and philosophers of science continue to debate the optimal placement of several elements where electronic configuration and chemical behavior diverge, as well as the visual geometry of the table itself.

The placement of the first two elements, hydrogen and helium, presents persistent classification ambiguities. Hydrogen ($1s^1$) has one valence electron like the Group 1 alkali metals, and under extreme pressures it can exhibit metallic properties; consequently, most tables place it atop Group 1. However, as a diatomic gas that can gain one electron to form a hydride ion ($H^-$) with a closed $1s^2$ shell, it also shares characteristics with the Group 17 halogens, while some tables float hydrogen separately above the grid as a standalone element. Helium ($1s^2$) possesses two s-orbital electrons, which strictly by electron configuration aligns it with the Group 2 alkaline earth metals in the s-block. Because its $1s$ shell is completely full, however, helium is chemically inert and is almost universally placed atop Group 18 with the p-block noble gases.

Another long-standing controversy concerns the composition of Group 3 in the 18-column table.³¹ Every standard table places scandium ($Z=21$) and yttrium ($Z=39$) in the top two spaces of Group 3, but published tables differ on which elements occupy the two spaces directly beneath them. One convention places lanthanum ($Z=57$) and actinium ($Z=89$) under yttrium, based on early spectroscopic assignments showing a single $d$ electron in their ground states. A competing convention places lutetium ($Z=71$) and lawrencium ($Z=103$) under yttrium, arguing that both possess filled $f$ subshells ($4f^{14}$ and $5f^{14}$), exhibit closer periodic trends to scandium and yttrium, and preserve an unbroken sequence of atomic numbers without splitting the d-block in a 32-column expansion. A third compromise leaves the two spaces blank with footnotes directing the reader to the 15-element lanthanoid and actinoid rows below.³² In 2015, IUPAC established a task group chaired by philosopher of chemistry Eric Scerri to examine the issue. In January 2021, the group published a provisional report recommending the scandium-yttrium-lutetium-lawrencium arrangement, though IUPAC has maintained that it does not mandate a single official periodic table layout.³³

Beyond the conventional rectangular grid, researchers have proposed more than 1,000 alternative two- and three-dimensional periodic visualizations to emphasize continuity or quantum symmetry. Prominent among these is the Left-Step Periodic Table, devised by French engineer Charles Janet in 1928.³⁴ Janet's system organizes elements strictly by the $n + \ell$ orbital-filling order, moving the s-block (including helium atop beryllium) to the far right edge of the chart and placing the f-, d-, and p-blocks in an unbroken ascending staircase. Another widely cited design is German-American chemist Theodor Benfey's 1964 two-dimensional spiral, which winds the elements in a continuous ribbon of increasing atomic number outward from hydrogen at the center, depicting the transition metals and inner transition series as protruding peninsulas.

All elements beyond uranium ($Z=92$), except for trace natural quantities of neptunium and plutonium, do not occur in nature and must be created artificially through nuclear fusion reactions in particle accelerators. The seventh period of the periodic table was officially completed in 2016 when IUPAC verified and named four superheavy elements: nihonium ($Z=113$), moscovium ($Z=115$), tennessine ($Z=117$), and oganesson ($Z=118$).

Experimental nuclear physicists at major facilities—including the Joint Institute for Nuclear Research (JINR) in Dubna, Russia; the RIKEN Nishina Center in Wako, Japan; the GSI Helmholtz Centre for Heavy Ion Research in Darmstadt, Germany; and Lawrence Berkeley National Laboratory (Berkeley Lab) in California—are working to synthesize element 119 (provisional systematic name ununennium) and element 120 (unbinilium).³⁵³⁶ The creation of either element will inaugurate the eighth period of the periodic table.

Synthesizing elements 119 and 120 requires a departure from the "hot fusion" method used to create elements 114 through 118, which relied on bombarding actinoid targets with a beam of neutron-rich calcium-48 ($Z=20$) ions.³⁷ Because target materials heavier than californium ($Z=98$) cannot be produced in sufficient microgram quantities to form an accelerator target, fusing calcium-48 cannot reach $Z=119$ or $Z=120$. Instead, researchers must accelerate heavier projectile beams, such as titanium-50 ($Z=22$) or vanadium-51 ($Z=23$), into targets of berkelium-249 or californium-249. In July 2024, scientists at Berkeley Lab's 88-Inch Cyclotron demonstrated a proof of concept for this pathway by bombarding plutonium-244 with a titanium-50 ion beam for 22 days, successfully producing two atoms of livermorium ($Z=116$) and validating the technique for dedicated searches for element 120.³⁸³⁹ Searches are also motivated by theoretical nuclear models predicting an "island of stability"—a region of superheavy nuclei around closed proton and neutron shells (such as neutron number $N=184$) where half-lives could increase from milliseconds to minutes or longer.

At the extreme end of the periodic table, relativistic effects significantly alter atomic orbitals and threaten to break down traditional periodic law.⁴⁰ In atoms with very high nuclear charges, inner-shell $s$ and $p$ electrons are accelerated to velocities approaching the speed of light. According to special relativity, their relativistic mass increases, causing the inner $s$ and $p$ orbitals to contract toward the nucleus and more effectively screen the nuclear charge from outer $d$ and $f$ orbitals, which consequently expand and destabilize. These relativistic shifts already explain why gold ($Z=79$) is yellow rather than silvery and why mercury ($Z=80$) is a liquid at room temperature. In the late seventh and eighth periods, relativistic effects are predicted to cause oganesson ($Z=118$) to behave as a reactive solid semiconductor rather than an inert noble gas, and to introduce an entirely new g-block of elements beginning around atomic number 121 where the $5g$ subshell starts to fill, blurring the distinct group periodicities seen in lighter rows.