Periodic Table and Periodic Trends – Vocabulary Flashcards

Vocabulary flashcards covering key terms and definitions from the lecture on classification of elements and periodic trends.

Periodic Law (Modern)

Statement that the physical and chemical properties of the elements are periodic functions of their atomic numbers.

Periodic Table (Long Form)

Modern arrangement of elements in seven periods and eighteen groups ordered by increasing atomic number, reflecting recurring properties.

Group

Vertical column in the periodic table; elements in the same group share similar valence-shell electronic configurations and chemical behaviour.

Period

Horizontal row in the periodic table; its number equals the highest principal quantum number (n) occupied by electrons of elements in that row.

s-Block Elements

Group 1 and 2 elements with outer electronic configuration ns¹–ns²; highly reactive, electropositive metals.

p-Block Elements

Elements of Groups 13–18 whose outer configuration ends in ns²np¹–ns²np⁶; include metals, metalloids and non-metals.

d-Block Elements

Groups 3–12 elements characterized by progressive filling of (n–1)d orbitals; commonly called transition metals.

f-Block Elements

Lanthanoids and actinoids in which (n–2)f orbitals are being filled; placed separately at the bottom of the periodic table.

Transition Elements

d-block metals that form coloured ions, exhibit variable oxidation states and often act as catalysts.

Inner-Transition Elements

f-block elements (lanthanoids & actinoids) with similar chemistry arising from f-orbital filling.

Metallic Character

Tendency of an element to lose electrons and form cations; increases down a group and decreases across a period.

Non-metallic Character

Tendency of an element to gain electrons and form anions; decreases down a group and increases across a period.

Atomic Radius

Half the distance between two identical atoms in a covalent molecule (non-metals) or metallic lattice (metals).

Ionic Radius

Effective size of a cation or anion measured from inter-ionic distances in ionic crystals; cations shrink, anions expand relative to parent atoms.

Ionization Enthalpy (First)

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

Second Ionization Enthalpy

Energy needed to remove a second electron from a gaseous mono-positive ion; always larger than the first.

Electron Gain Enthalpy

Enthalpy change when a gaseous atom accepts an electron to form an anion; negative values indicate energy release.

Electron Affinity

Numerical value equal in magnitude but opposite in sign to electron gain enthalpy, often expressed as a positive quantity when energy is released.

Electronegativity

Relative tendency of an atom in a compound to attract shared electrons; not directly measurable but expressed on scales such as Pauling.

Pauling Scale

Most widely used numerical scale of electronegativity, assigning fluorine the maximum value of 4.0.

Valence

Combining capacity of an element, often equal to the number of outer-shell electrons or eight minus that number for main-group elements.

Oxidation State

Apparent charge an atom acquires in a compound based on electronegativity considerations; modern term replacing valence in many contexts.

Diagonal Relationship

Similarity in properties between the first element of a group and the second element of the next group (e.g., Li with Mg, Be with Al).

Amphoteric Oxide

Oxide that behaves as an acid with bases and as a base with acids (e.g., Al₂O₃, ZnO).

Basic Oxide

Oxide that reacts with water to give a base; typical of metals on the left of a period (e.g., Na₂O → 2 NaOH).

Acidic Oxide

Oxide that reacts with water to give an acid; typical of non-metals on the right of a period (e.g., Cl₂O₇ → 2 HClO₄).

Neutral Oxide

Oxide exhibiting neither acidic nor basic behaviour, such as CO or N₂O.

Periodic Trends

Systematic variations in properties (atomic radius, ionization enthalpy, etc.) observed across periods or down groups.

Noble-Gas Configuration

Stable valence-shell arrangement ns²np⁶ (or 1s² for He) that many atoms achieve through ion formation or bonding.

IUPAC Nomenclature (Z > 100)

System using Latin roots for digits (nil, un, bi, tri…) followed by –ium to assign temporary names and symbols to super-heavy elements.

Isoelectronic Species

Atoms or ions having identical numbers of electrons (e.g., O²⁻, F⁻, Na⁺, Mg²⁺ all contain 10 electrons).

Shielding (Screening) Effect

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

Effective Nuclear Charge (Zₑff)

Net positive charge experienced by a valence electron after accounting for shielding; increases across a period.

Dobereiner’s Triads

Early classification grouping sets of three elements with similar properties where the atomic weight of the middle element is the mean of the other two.

Law of Octaves

Newlands’ observation that every eighth element shows similar properties when elements are arranged by increasing atomic weight (valid up to Ca).

Mendeleev’s Periodic Table

1869 arrangement of elements by increasing atomic weight, leaving gaps for undiscovered elements and correcting masses to keep similar elements together.

Moseley’s X-ray Experiment

1913 study showing that square-root of X-ray frequency varies linearly with atomic number, establishing Z as the fundamental basis for periodicity.

Actinoid Series

Row of 14 inner-transition elements from Th (90) to Lr (103) involving 5f orbital filling; predominantly radioactive.

Lanthanoid Series

Row of 14 inner-transition elements from Ce (58) to Lu (71) involving 4f orbital filling; exhibit lanthanoid contraction.

Aufbau Principle

Rule that electrons occupy available orbitals in order of increasing energy, underlying the construction of electronic configurations.