AVS Ki Pathshala • Elements Encyclopedia
Systematic names • research/learning extension • not official named elements
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AVS Ki Pathshala • Complete reference for students
Döbereiner noticed that some chemically similar elements could be arranged in groups of three. The atomic mass of the middle element was approximately the arithmetic mean of the other two. Examples: Li–Na–K and Ca–Sr–Ba. Why it mattered: it was one of the first recognisable attempts to connect atomic mass with repeating chemical behaviour.
Alexandre-Émile Béguyer de Chancourtois arranged elements by atomic weight on a cylindrical spiral. Similar elements appeared at recurring vertical positions. It was an early visual demonstration of periodicity, although the presentation was not widely understood at the time.
John Newlands arranged elements by increasing atomic weight and observed that chemical similarities often reappeared at roughly every eighth position among the lighter elements. The analogy with musical octaves earned the name “Law of Octaves”. Its main limitation was that it did not accommodate all known elements cleanly and offered no convincing place for undiscovered elements.
Lothar Meyer independently demonstrated periodic relationships between atomic mass and physical properties such as atomic volume. His work strongly supported the idea that element properties repeat in a systematic manner.
Dmitri Mendeleev arranged elements primarily by increasing atomic mass while giving priority to recurring chemical properties. Crucially, he left gaps for elements that had not yet been discovered and predicted their properties. The later discoveries of gallium, scandium and germanium provided famous confirmations of those predictions.
Core insight: a periodic table is not just a list — it is a predictive model.
Mendeleev used prefixes such as eka- to describe elements expected below known elements. Eka-aluminium became gallium; eka-boron became scandium; and eka-silicon became germanium. Their properties were sufficiently close to his predictions to make the periodic law much more convincing.
The discovery of helium, neon, argon, krypton, xenon and radon showed that an entire family of chemically distinctive elements had to be incorporated into the classification. This led to the modern Group 18 position.
Henry Moseley used X-ray spectra to establish a systematic relationship between X-ray frequency and atomic number. This showed that atomic number, not atomic mass, is the correct fundamental ordering principle. Several apparent anomalies in mass-based tables were thereby resolved.
The physical and chemical properties of elements are periodic functions of their atomic numbers. Modern periodicity is explained by repeating patterns of electron configurations, especially valence-shell configurations.
Principal shells, subshells, orbitals and quantum numbers provide the electronic foundation of the modern table. Aufbau filling, the Pauli exclusion principle and Hund's rule explain why the s-, p-, d- and f-blocks occupy their characteristic positions.
After uranium, nuclear reactions and particle accelerators opened the route to elements beyond the naturally occurring heavy elements. Many superheavy elements are synthetic and exist only for very short times, making their discovery a triumph of nuclear chemistry and physics.
IUPAC temporary systematic names are constructed from numerical roots when a new element has not yet received an approved permanent name. Once discovery is accepted, the discoverers can propose a name, followed by IUPAC review and approval. Temporary systematic symbols are distinct from permanent official symbols.
Oganesson (Z = 118) is the latest officially named element. Z = 119 onward remains a research frontier. The 119–127 extension on this page is therefore a learning/research layer, not a claim that those elements have been experimentally discovered or permanently named.
Across a period: generally decreases because nuclear charge increases while electrons are added to the same principal shell. Down a group: increases because new shells are added. Trend is affected by shielding, subshell structure and unusual bonding definitions of atomic radius.
Cations are generally smaller than their parent atoms because electron loss reduces electron–electron repulsion and may remove a shell. Anions are generally larger because extra electrons increase repulsion. In an isoelectronic series, radius generally decreases as nuclear charge increases.
Ionization energy is the energy required to remove an electron from a gaseous atom/ion. It generally increases across a period and decreases down a group. Common exceptions arise from subshell energies and electron pairing, such as Be→B and N→O comparisons.
IE₁ < IE₂ < IE₃ … because each successive electron is removed from an increasingly positive ion. A very large jump indicates that the next electron belongs to an inner shell. This is useful for predicting valency and group position.
It describes the energy change associated with adding an electron to a gaseous atom. Halogens strongly favour electron gain. The trend is less smooth than ionization energy because of electron–electron repulsion, subshell stability and atomic size. Noble gases generally resist addition of an electron.
Electronegativity is the tendency of an atom in a chemical bond to attract shared electron density. It generally increases across a period and decreases down a group. Fluorine has the highest Pauling electronegativity.
The outer electron experiences the nucleus through the screening of other electrons. A simple conceptual relation is Zeff ≈ Z − S, where S represents shielding. Across a period, Z rises while shielding does not rise equally, so effective attraction generally increases.
Inner-shell electrons partially shield outer electrons from the full nuclear charge. Shielding increases down a group because more inner shells are present. In the same shell, shielding is not equally effective for all subshells.
Metallic character generally increases down a group and decreases across a period. Metals more readily lose electrons and form cations. Non-metallic character broadly follows the reverse trend.
For many main-group elements, common valency is linked to the number of valence electrons and the tendency to attain a stable configuration. Transition elements are different because both ns and (n−1)d electrons can participate, producing variable oxidation states.
Main-group oxidation states often follow predictable valence patterns, while d-block elements commonly show several oxidation states. The stability of oxidation states changes across periods and down groups because orbital energies and bonding environments change.
Across a period, oxides often progress broadly from basic/ionic → amphoteric → acidic/covalent. Down a group, metallic character usually increases, often strengthening the basic character of corresponding oxides.
For alkali metals, reactivity generally increases down the group because electron removal becomes easier. For many metals, reducing character follows related patterns, but actual reactivity also depends on lattice energy, hydration and reaction conditions.
Halogen oxidising ability generally decreases down Group 17 because the larger atoms attract an incoming electron less strongly. Fluorine is exceptionally strong as an oxidising agent.
Increasing electronegativity difference tends to increase bond polarity. Small, highly charged cations can strongly polarise large anions, producing covalent character — a useful connection between periodic position and bonding.
There is no single universal left-to-right or top-to-bottom rule. Values depend on bonding and structure: metallic bonding, giant covalent networks, molecular forces and atomic mass all matter. Carbon and tungsten illustrate why structural bonding can dominate simple mass trends.
Density often changes irregularly because both mass and atomic/structural volume change. Therefore density should be treated as a property with periodic tendencies in families, not as a perfectly monotonic periodic trend.
Metals generally conduct electricity because of mobile electrons. Non-metals are usually poor conductors, while metalloids and semiconductors such as silicon occupy an important intermediate regime. Conductivity is therefore closely tied to electronic structure.
Electrons in the outermost occupied shell that commonly participate in bonding and determine many chemical properties.
The combining capacity of an element, often related to the number of electrons lost, gained or shared to achieve a more stable electronic arrangement.
Reduction in the effective nuclear attraction on an outer electron caused by repulsion from inner electrons.
The net positive attraction experienced by an electron after the shielding effect of other electrons is considered.
The extent to which an orbital electron density approaches the nucleus. For orbitals in the same principal shell, penetration generally follows s > p > d > f.
Similarities between diagonally placed elements, especially in the second and third periods, caused by comparable size, charge density and polarising power. Classic examples include Li–Mg and Be–Al.
In heavier p-block elements, the outer ns² pair may participate less readily in bonding. This can stabilise lower oxidation states, especially toward the bottom of groups 13–16.
The gradual decrease in atomic and ionic radii across the lanthanides caused mainly by the relatively poor shielding of nuclear charge by 4f electrons.
Atoms or ions having the same number of electrons. In an isoelectronic series, the species with greater nuclear charge is generally smaller.
Electrons occupy available orbitals in an order of increasing energy. The common filling sequence explains the broad block structure of the periodic table.
No two electrons in an atom can have the same complete set of four quantum numbers. An orbital therefore accommodates a maximum of two electrons with opposite spins.
Degenerate orbitals are singly occupied with parallel spins before pairing occurs. This lowers repulsion and helps explain ground-state configurations.
Shielding describes how other electrons reduce nuclear attraction; penetration describes how strongly an electron density approaches the nucleus. Together they help explain subshell energy order and periodic exceptions.
A measure of atomic size. Different experimental definitions exist, such as covalent, metallic and van der Waals radii, so quoted values depend on the bonding context.
Energy needed to remove an electron from a gaseous species. First ionization energy concerns a neutral atom; successive ionization energies concern increasingly positive ions.
A term used for the energetic change when an electron is added to a gaseous atom. Sign conventions vary between sources, so always check whether the source reports affinity or electron-gain enthalpy.
A relative scale describing how strongly an atom in a bond attracts shared electron density. Pauling, Mulliken and Allred–Rochow are different electronegativity scales.
The number of protons in the nucleus. It defines the identity of an element and is the fundamental ordering variable of the modern periodic table.
Total number of protons and neutrons in a particular nuclide: A = Z + N. Isotopes of an element have the same Z but different A.
Atoms of the same element with the same atomic number but different numbers of neutrons. Their chemical behaviour is broadly similar, while nuclear stability and mass can differ greatly.
Nuclides having the same mass number but different atomic numbers. They therefore belong to different elements.
Nuclides having the same number of neutrons but different atomic numbers.
Repetition of comparable physical and chemical properties at regular intervals when elements are arranged by increasing atomic number.
The s-, p-, d- and f-block labels indicate the subshell into which the differentiating electron enters in the ground-state configuration.
A group is a vertical column; a period is a horizontal row. Elements in a group often share valence-electron patterns and related chemistry, while periods reflect progressive filling of shells/subshells.
A periodic trend is a broad systematic variation across a row or column. A diagonal relationship is a local similarity between diagonally positioned elements and should not be confused with a general trend.
For not-yet-permanently-named superheavy elements, IUPAC numerical roots are combined to form temporary systematic names and symbols. These are placeholders, not evidence that the element has been experimentally confirmed.