Period Periodic Table

A period in the periodic table refers to a horizontal row of chemical elements. Each element across a given period contains the same number of electron shells, while the number of protons in the nucleus increases from left to right. This progressive increase in atomic number results in a gradual reduction in metallic character across the row. The alignment of elements in vertical groups reflects their broadly similar chemical and physical properties, a relationship explained by the periodic law. For instance, the halogens, found in the penultimate group, share reactivity patterns and exhibit a characteristic tendency to gain one electron to achieve a noble gas configuration. Presently, 118 elements have been identified and placed within seven complete periods.

Underlying Principles of Periodicity

The arrangement of orbitals by increasing energy is described through the Madelung rule, in which the order of orbital filling corresponds to diagonal sets of quantum numbers. Modern quantum mechanics explains periodic trends such as atomic radius, ionisation energy and electronegativity by the structured filling of electron shells. As the atomic number ascends, electrons populate orbitals in the order predicted by the rule. Completion of a shell corresponds with the end of a period.
Patterns across periods vary depending on the block of the periodic table. Elements in the f-block and p-block often show limited horizontal similarity, with vertical relationships down groups proving more significant. However, the d-block exhibits pronounced across-period trends, while f-block elements display a marked uniformity across periods due to similarities in electron configuration.

Overview of Periods

There are currently seven completed periods. Any future synthesis of additional elements would extend the table into an eighth period, potentially involving a new g-block. Not all predicted elements are experimentally attainable, raising uncertainty regarding the existence of a ninth period.

Period 1

Period 1 contains only two elements: hydrogen and helium, the fewest of any period. These atoms follow a duplet rule rather than the octet rule, as they occupy only the 1s orbital.
Hydrogen displays chemical versatility, behaving like a group 1 element when donating an electron and like a group 17 element when gaining one. It is the most abundant element in the universe and forms the bulk of stellar matter in the plasma state. On Earth, molecular hydrogen is relatively rare and is commonly produced industrially from methane.
Helium exists primarily as a gas and is the second most abundant element in the universe. Most helium originated during the Big Bang, with additional quantities produced through nuclear fusion in stars. On Earth it is comparatively scarce, occurring mainly as a product of radioactive decay trapped within natural gas reserves.

Period 2

Period 2 contains the s- and p-block elements: lithium, beryllium, boron, carbon, nitrogen, oxygen, fluorine and neon. These include several biologically essential elements.

• Lithium, the lightest metal, is extremely reactive and is found only in compounds.
• Beryllium possesses a high melting point. Though formed in minute amounts during the Big Bang, it was largely consumed in stellar nucleosynthesis.
• Boron occurs in nature as borates and is an essential micronutrient for plant development.
• Carbon, the fourth most abundant element in the universe, forms the structural basis of organic chemistry.
• Nitrogen, comprising 78% of Earth’s atmosphere, is an essential component of amino acids and nucleic acids.
• Oxygen, making up 21% of the atmosphere and a major component of the crust, is crucial for respiration.
• Fluorine, the most reactive non-metal, is never found free in nature.
• Neon is a noble gas widely used in lighting.

Period 3

All period 3 elements occur naturally and possess at least one stable isotope. With the exception of argon, all are integral to geology or biology.

• Sodium, abundant in seawater, is essential for fluid regulation.
• Magnesium, found in chlorophyll, is needed for photosynthesis.
• Aluminium, the most abundant metal in Earth’s crust, is a lightweight structural material.
• Silicon, a metalloid, is vital to the semiconductor industry and forms the basis of silicate minerals.
• Phosphorus is essential to DNA and energy transfer but is never found free due to high reactivity.
• Sulphur, present in amino acids, is crucial for protein structure.
• Chlorine forms salts such as sodium chloride and is widely used as a disinfectant.
• Argon, a noble gas, is used in lighting to protect hot filaments.

Period 4

Period 4 introduces the first d-block transition metals. It includes biologically significant elements such as potassium and calcium, essential for nerve transmission and bone structure. Transition metals present in this period include scandium, titanium, vanadium, chromium, manganese, iron, cobalt, nickel and copper. Iron is the heaviest element synthesised in large quantities in stars and is a major component of Earth’s core. The period concludes with six p-block elements: gallium, germanium, arsenic, selenium, bromine and krypton.

Period 5

Period 5 mirrors the structure of period 4, containing transition metals and p-block elements. This period includes two biologically important heavy elements — molybdenum and iodine — as well as technetium, the lightest element with no stable isotopes. It contains one additional post-transition metal and one fewer non-metal than the previous period.

Period 6

Period 6 is the first to incorporate the f-block, consisting of the lanthanide series. These rare earth elements exhibit notable similarities due to gradual filling of 4f orbitals. The period also contains the heaviest stable elements. While many heavy metals in this row are toxic or radioactive, elements such as platinum and gold are chemically inert and valued for catalytic and ornamental applications.

Period 7

All elements in period 7 are radioactive. It includes plutonium, the heaviest naturally occurring element, followed by a sequence of primarily synthetic elements. Some, such as americium and curium, can be produced in measurable quantities, whereas others exist only fleetingly in laboratory conditions.
Chemical trends in this period are less consistent than in earlier rows due to factors such as spin–orbit coupling and relativistic effects, arising from the extremely high nuclear charge of these superheavy atoms. The actinides demonstrate particularly diverse oxidation states and behaviours.

Period 8

An eighth period has been theorised and may include elements belonging to a g-block, though none have yet been synthesised. Theoretical models suggest physical limits may prevent the existence of all predicted superheavy elements, and the structure of such a period remains uncertain.