Periodic table

The periodic table of the elements is a systematic and ordered arrangement of all known chemical elements, organised into horizontal periods and vertical groups according to recurring patterns in their physical and chemical properties. It serves as a cultural and scientific icon in chemistry and is equally fundamental to physics, materials science, and other natural sciences. Its modern layout embodies the principles of periodicity, illustrating how elemental behaviour varies in relation to atomic number and electron configuration.

Historical Development and Periodic Law

The earliest generally accepted form of the periodic table was developed by the Russian chemist Dmitri Mendeleev in 1869. Working at a time when not all elements were known, Mendeleev arranged the elements in order of increasing atomic mass and noticed repeating patterns of properties. This insight led him to formulate the periodic law, stating that the chemical properties of elements recur periodically when arranged by atomic mass. Remarkably, he left intentional gaps for then-undiscovered elements and successfully predicted the existence and properties of several of these, thereby validating the periodic law.
A more complete theoretical foundation emerged in the early twentieth century with the discovery of the atomic number, which was shown to represent the number of protons in an atom’s nucleus. Henry Moseley’s work established that atomic number, not atomic mass, was the correct basis for ordering the elements. Concurrent advances in quantum mechanics provided further insights by elucidating the structure of electron shells and subshells, offering an explanation for the observed periodicity in elemental behaviour.
A recognisably modern form of the periodic table was achieved in 1945 after Glenn T. Seaborg identified that the actinides form an f-block series rather than belonging within the d-block. This transformation clarified the internal arrangement of the table and provided the foundation for its contemporary structure.

Structure and Organisation of the Periodic Table

The standard periodic table lists elements in order of increasing atomic number. A new period begins when a new electron shell starts to fill, while groups bring together elements with similar outer electron configurations. In general, elements belonging to the same group exhibit comparable chemical behaviour. For instance, oxygen, sulphur, and selenium all possess four electrons in their outermost p-subshells, leading to broadly similar chemical characteristics.
The table is divided into four blocks based on the subshell in which the differentiating electron resides:

• s-block, comprising groups 1 and 2;
• p-block, containing groups 13 to 18;
• d-block, also known as the transition metals;
• f-block, containing the lanthanides and actinides.

The f-block elements, though logically positioned between the alkaline earth metals and the scandium group, are conventionally displayed below the main table to preserve horizontal compactness. A traditional diagonal line is often drawn to separate metals and non-metals, reinforcing their broad differences in physical properties, though many borderline cases exist.

Periodic Trends

The periodic table visually represents several recurring chemical trends:

• Metallic character increases down a group and from right to left across a period.
• Non-metallic character strengthens from the bottom left to the top right of the table.
• Atomic radius, ionisation energy, electron affinity, and electronegativity each follow characteristic patterns governed by nuclear charge and electron shielding.

These predictable variations allow chemists to make informed predictions about reactivity, types of bonding, oxidation states, and other fundamental behaviours.

Chemical Elements, Isotopes, and Atomic Structure

Each chemical element is defined by a unique atomic number, commonly represented by its chemical symbol. Hydrogen (H) has atomic number 1, helium (He) 2, and lithium (Li) 3, progressing sequentially to the heaviest known elements.
While the number of protons determines the identity of an element, the number of neutrons may vary. Atoms with the same number of protons but different neutron counts are known as isotopes. Naturally occurring elements often comprise mixtures of isotopes, each contributing to the element’s overall atomic weight. For elements with no stable or abundant natural isotopes, the mass of the most stable artificial isotope is used.
Isotopes are always grouped together under a single element within the periodic table, as they share identical electronic structures and chemical behaviour.

Naturally Occurring and Synthetic Elements

At present, 118 elements are known, completing the first seven periods. Of these, the first 94 exist naturally on Earth. Uranium, with atomic number 94, can undergo spontaneous fission, enabling the formation of minute quantities of neptunium and plutonium through neutron capture and subsequent beta decay.
Some elements heavier than uranium have been detected in exceptional astrophysical environments. Elements up to einsteinium (atomic number 99) have been observed in the peculiar star Przybylski’s Star. Elements up to fermium (100) are believed to have been generated in the natural fission reactor at Oklo in Gabon, though these have since decayed.
The heaviest elements, such as those from americium (95) to oganesson (118), have been synthesised exclusively in laboratories through advanced nuclear reactions. The last four elements were formally named in 2016, completing the seventh period. However, detailed chemical characterisation of the superheavy elements remains ongoing, as their extreme instability makes experimental study challenging.
Supernovae and neutron-star mergers may create even heavier elements via the r-process, though this has not yet been experimentally confirmed. Theoretical calculations suggest possible formation of nuclides with mass numbers between 280 and 290, but most are predicted to decay in extremely short timescales.

Group Names, Numbering, and Classification

Modern international convention numbers the groups from 1 to 18, starting with the alkali metals and concluding with the noble gases. This system, established by IUPAC in 1988, replaced earlier naming systems that used Roman numerals and varied between countries. Some groups are alternatively referred to by the name of their first element, such as the scandium group for group 3. The f-block is not assigned group numbers under the current scheme.
Earlier classifications featured the s- and p-block groups labelled with suffix A, and the d-block groups with suffix B, though such nomenclature is now deprecated. Historically, groups 8, 9, and 10 were sometimes treated collectively as group VIII.

Presentation Forms of the Periodic Table

The periodic table exists in several forms, though the 18-column and 32-column layouts are the most widely recognised. The 32-column “long form” table includes the f-block in its natural position between groups 2 and 3. However, for practicality, most printed tables display the lanthanides and actinides beneath the main body.
Many alternative periodic table designs have been proposed, including spiral, circular, and left-step arrangements. These seek to emphasise different aspects of periodicity, such as electron configurations or elemental relationships. Nevertheless, the standard form remains the most universally adopted due to its clarity and compatibility with chemical education.

Contemporary Developments and Future Prospects

The periodic table continues to evolve alongside scientific advances. As experimental methods in nuclear physics improve, it is expected that new synthetic elements may be produced, expanding the table into the so-called extended periodic table. However, theoretical calculations indicate that the behaviour of these hypothetical superheavy elements may diverge significantly from established periodic patterns due to relativistic effects on electron orbitals.