Periodic Table and Periodicity – Core Vocabulary

A set of key vocabulary flashcards covering historical development, structural features, periodic trends, and fundamental concepts of the periodic table and periodicity.

Periodic Table

A tabulated classification of chemical elements arranged by increasing atomic number, showing periodic repetition of physical and chemical properties.

Periodicity

The recurring trends in element properties when arranged by increasing atomic number.

Period

A horizontal row in the periodic table in which elements have the same number of electron shells.

Group

A vertical column of the periodic table whose elements share similar valence-shell electron configurations and properties.

Prout’s Hypothesis

Early proposal that atomic masses are simple whole-number multiples of hydrogen’s mass.

Dobereiner’s Triads

Sets of three elements where the middle element’s atomic mass is the arithmetic mean of the other two.

Newland’s Law of Octaves

Observation that every eighth element, when arranged by increasing atomic mass, had similar properties.

Lothar Meyer’s Volume Curve

Graph plotting atomic volume versus atomic mass, revealing periodic maxima and minima for element families.

Mendeleev’s Periodic Law

Statement that element properties are periodic functions of their atomic masses (later superseded).

Modern Periodic Law

Principle that element properties are periodic functions of their atomic numbers.

s-Block Elements

Elements in groups 1 and 2 (plus He) whose valence electrons enter s orbitals.

p-Block Elements

Elements in groups 13–18 whose valence electrons enter p orbitals.

d-Block Elements

Transition elements in which valence electrons enter (n–1)d orbitals.

f-Block Elements

Inner–transition elements (lanthanides and actinides) where valence electrons enter (n–2)f orbitals.

Alkali Metals

Group-1 elements characterized by ns¹ outer configuration and strong reactivity.

Alkaline Earth Metals

Group-2 elements with ns² outer configuration and +2 valence.

Transition Elements

d-block metals exhibiting variable oxidation states, colored ions, and catalytic activity.

Inner-Transition Elements

f-block metals (lanthanides & actinides) with progressive filling of f subshells.

Lanthanide Contraction

Steady decrease in atomic/ionic radii from La to Lu caused by poor shielding of 4f electrons.

Effective Nuclear Charge (Z*)

Net positive charge experienced by an electron after accounting for shielding by inner electrons.

Shielding Effect

Reduction in nuclear attraction on valence electrons due to repulsion by inner-shell electrons.

Slater’s Rules

Empirical guidelines for estimating shielding constants to calculate effective nuclear charge.

Covalent Radius

Half the distance between nuclei of two identical atoms joined by a single covalent bond.

Van der Waals Radius

Half the closest approach distance between two non-bonded atoms of the same element.

Metallic Radius

Half the internuclear distance between adjacent atoms in a metallic crystal lattice.

Ionic Radius

Effective distance from nucleus to the outermost electron in a cation or anion.

Isoelectronic Species

Atoms or ions possessing identical electron counts and configurations.

Ionization Energy (IE)

Energy required to remove the most loosely bound electron from a gaseous atom.

Successive Ionization Energies

Energies needed to remove second, third, etc., electrons; each is higher than the previous.

Electron Affinity (EA)

Energy released when a gaseous atom gains an electron to form an anion.

Electronegativity

Relative tendency of a bonded atom to attract shared electrons toward itself.

Pauling Scale

Numerical scale (0–4) quantifying electronegativity based on bond-energy differences.

Mulliken Electronegativity

Average of an element’s ionization energy and electron affinity expressed in electron volts.

Allred–Rochow Scale

Electronegativity derived from effective nuclear charge divided by covalent radius squared.

Sanderson Electronegativity

Scale using stability ratio (electron density relative to noble gas) to express electronegativity.

Atomic Radius Trend

Decreases across a period and increases down a group due to nuclear charge and shell addition.

Ionization Energy Trend

Generally rises across a period and falls down a group, with exceptions for stable sub-shells.

Electron Affinity Trend

Becomes more negative across a period and less negative down a group; Cl is most exothermic.

Electronegativity Trend

Increases from left to right across periods and decreases down groups; F is highest, Cs lowest.

Oxidizing Power Trend

Increases across a period and decreases down a group among non-metals (F₂ > Cl₂ > Br₂ > I₂).

Reducing Power Trend

Decreases across a period and increases down a group among metals.

Diagonal Relationship

Similarity between certain diagonally adjacent pairs (e.g., Li-Mg, Be-Al) in the periodic table.

Half-Filled & Fully-Filled Stability

Extra stability associated with exactly half-filled or completely filled subshells (e.g., p³, p⁶).

Numerical Roots (IUPAC)

Prefixes (nil, un, bi, tri, quad, pent, hex, sept, oct, enn) used for systematic naming of elements Z>100.

Copernicium (Cn)

Element 112; predicted as ‘eka-mercury’ before discovery.

Blocks of Periodic Table

Regions (s, p, d, f) classified by the subshell receiving the last electron in ground-state atoms.

Period Capacity

Maximum element counts per period: 2, 8, 8, 18, 18, 32, 32 for periods 1–7 respectively.

Shielding Constant (S)

Numerical value representing total screening effect of inner electrons in Slater’s formula.

Atomic Volume

Volume occupied by one gram-atom of an element in solid state; shows periodic variation.

Percentage Ionic Character

Degree of ionicity in a bond estimated from electronegativity difference (Hanny–Smith relation).

Basic Oxide

Oxide that reacts with acids producing salt and water; formed by highly electropositive metals.

Acidic Oxide

Oxide that reacts with bases; typical of non-metals with high electronegativity.

Amphoteric Oxide

Oxide capable of reacting as both acid and base, e.g., Al₂O₃.

Naming of Hydrides

Less electronegative element named first, followed by the more electronegative plus suffix ‘-ide’ (e.g., hydrogen fluoride).

Electron Configuration

Arrangement of electrons in atomic orbitals, used to deduce period, group, and block placement.

Lanthanides

14 elements from Ce (58) to Lu (71) with progressive 4f filling, often called rare earths.

Actinides

14 elements from Th (90) to Lr (103) with progressive 5f filling; many are radioactive.

Shielding (Screening) Effect

Repulsion by inner electrons that diminishes the nucleus’ pull on outer electrons, lowering Z*.

Penetration Power

Relative closeness of orbital types to the nucleus (s > p > d > f) influencing ionization energy.

Successive Electron Affinities

Energy changes for adding second, third electrons; always endothermic due to anion–electron repulsion.