Explore electron configurations, periodic trends, side-by-side element comparisons, interactive Bohr models, and chemistry quizzes.
The first Bose-Einstein condensate in history was produced using Rubidium-87 atoms cooled to 170 nanokelvin by Cornell and Wieman in 1995.
The Periodic Table of Chemical Elements is the master tabular arrangement of all known chemical elements, organized in increasing order of their atomic number (Z), electron configurations, and recurring chemical properties. It is widely considered one of the most significant achievements in modern science, establishing a universal framework that unifies physics, chemistry, geology, and biology.
In 1869, Russian chemist Dmitri Mendeleev formulated the first widely accepted periodic system by organizing the 63 then-known elements by atomic mass. Mendeleev famously left deliberate blank spaces in his chart for undiscovered elements (such as Gallium, Germanium, and Scandium) and accurately predicted their physical and chemical properties decades before their discovery.
In 1913, English physicist Henry Moseley discovered using X-ray spectroscopy that the true fundamental property of each element is its nuclear charge (atomic number), rather than its atomic mass. This gave rise to the Modern Periodic Law:
The modern IUPAC periodic table is organized systematically into four hierarchical dimensions:
Every element is sequenced incrementally by nuclear charge from Hydrogen (Z = 1) to Oganesson (Z = 118). In neutral atoms, atomic number also denotes total electron count.
18 vertical columns. Elements within the same group share identical valence shell configurations, yielding matching valencies and related reaction pathways.
7 horizontal rows. The period index matches the principal quantum number (n = 1 to 7) of the outermost occupied shell in ground-state neutral atoms.
Elements are categorized by the orbital type into which the differentiating (last) electron enters according to the Aufbau principle.
Under IUPAC nomenclature, columns are numbered consecutively from 1 to 18. Here is the canonical breakdown critical for competitive exams:
Lithium to Francium (ns¹ valence). Extremely reactive, low melting points, soft, react vigorously with water to form basic hydroxides and hydrogen gas.
Beryllium to Radium (ns² valence). Form +2 cations; higher melting points and densities than Group 1. Basic oxides historically called 'earths'.
Incompletely filled d-subshells. Characterized by variable oxidation states, coloured ions/complexes, catalytic properties, and paramagnetism.
ns² np³ valence with a stable half-filled p subshell. Exhibits allotropy and oxidation states from -3 to +5.
ns² np⁴ valence. Derived from Greek 'chalcos' (copper/ore) + 'gen' (former); most metal ores exist as oxides or sulfides.
ns² np⁵ valence. Highly electronegative, requiring 1 electron for stable octet. React with alkali metals to yield salts like NaCl.
Helium (1s²) and Ne–Rn (ns² np⁶). Possess completely filled valence octets, imparting near-zero electron affinity, extraordinarily high first ionization potentials, and monatomic chemical inertness under standard terrestrial conditions.
Periodic trends are predictable patterns in elemental properties across groups and periods caused by the balance between increasing nuclear charge (Z_eff) and electron shielding:
Across a Period: Atomic radius decreases from left to right because additional protons increase effective nuclear charge (Z_eff), drawing valence electrons closer without adding extra shells.
Down a Group: Atomic radius increases as each new period adds an entire principal quantum shell (n), pushing valence electrons further from the nucleus.
Trend Direction: The measure of an atom's ability to attract shared electron pairs in a chemical bond. Fluorine is the most electronegative element (3.98 Pauling units), while Cesium and Francium are the least (~0.7).
Down a Group: Electronegativity decreases because valence bonding orbitals are located further from the positive nucleus with greater core shielding.
The minimum energy required to remove the most loosely bound electron from an isolated gaseous neutral atom. Helium possesses the highest ionization potential (2372 kJ/mol).
Reflects the tendency of an atom to lose electrons and form positive cations (electropositive character). Decreases from left to right across a period and increases down a group.
Direct, authoritative answers to common conceptual chemistry questions:
Groups are the 18 vertical columns in the periodic table. Elements in the same group possess identical valence electron configurations, leading to highly similar chemical properties and reactive behaviors (for example, Group 1 alkali metals all have 1 valence electron).
Periods are the 7 horizontal rows in the periodic table. Each period corresponds to the principal energy level (electron shell, n = 1 to 7) being filled by electrons across the elements in that row.
There are currently 118 officially recognized chemical elements, ranging from Hydrogen (atomic number 1) to Oganesson (atomic number 118). Elements 1 through 94 occur naturally on Earth, while elements 95 through 118 are synthetic elements synthesized in particle accelerators.
Lanthanides (elements 57–71) and Actinides (elements 89–103) represent the f-block where 4f and 5f electron subshells are progressively filled. If placed directly between Groups 2 and 3, the periodic table would be 32 columns wide, making standard desktop and textbook viewing awkward without adding new structural clarity.
Under Henry Moseley’s Modern Periodic Law, an element’s position is strictly determined by its atomic number (Z), which is the number of protons in the nucleus of its atom, and its ground-state electron configuration.
The atomic number (Z) is an exact integer representing the number of protons inside the atom’s nucleus. Atomic weight (relative atomic mass) is a weighted average of the atomic masses of all naturally occurring stable isotopes of that element, measured relative to Carbon-12.