Chemical Periodicity & Modern Periodic Table – Vocabulary Review

A comprehensive set of vocabulary flashcards covering key terms, laws and properties discussed in the lecture on Chemical Periodicity and the Modern Periodic Table.

Periodic Table

Tabular arrangement of elements in order of increasing atomic number so that elements with similar properties fall in the same columns (groups).

Dobereiner’s Law of Triads

1829 classification in which groups of three elements had similar properties and the atomic weight of the middle element was the arithmetic mean of the other two.

Telluric Screw (Chancourtois Helix)

1862 three-dimensional spiral listing of elements by atomic weight that demonstrated periodic recurrence of properties.

Newlands’ Law of Octaves

1864 ordering of elements by atomic weight showing that every eighth element had properties similar to the first, analogous to musical octaves.

Mendeleev’s Periodic Law

Statement that the physical and chemical properties of elements are periodic functions of their atomic weights; basis of Mendeleev’s 1869 table.

Modern Periodic Law

Moseley’s rule that the properties of elements are periodic functions of their atomic numbers.

Group (Periodic Table)

Vertical column of the periodic table; elements within a group share similar valence-shell electron configurations and properties.

Period (Periodic Table)

Horizontal row of the periodic table; elements show a gradation of properties as atomic number increases across a period.

s-Block Elements

Elements in Groups 1 and 2 whose valence electrons occupy s orbitals; also called alkali metals and alkaline-earth metals.

p-Block Elements

Elements in Groups 13–18 with valence electrons entering p orbitals; include metals, metalloids and non-metals.

d-Block Elements

Elements in Groups 3–12 in which the (n–1)d subshell is being filled; known as transition elements.

f-Block Elements

Lanthanides and actinides in which (n–2)f orbitals are being filled; called inner-transition elements.

Representative (Normal) Elements

s- and p-block elements whose outermost shell is incomplete while inner shells are complete; show wide range of chemical behavior.

Transition Elements

d-block metals with incomplete d subshells in atoms or common ions; exhibit variable oxidation states and colored compounds.

Inner-Transition Elements

f-block elements (lanthanides and actinides) characterized by filling of f orbitals and similar chemistry within each series.

Lanthanide Series

14 elements from Ce (58) to Lu (71) in which 4f orbitals are progressively filled; also called rare earths.

Actinide Series

14 elements from Th (90) to Lr (103) in which 5f orbitals are being filled; many are radioactive and include transuranic elements.

Transuranic Elements

Elements with atomic numbers greater than 92 (uranium); all are synthetic and radioactive.

Noble Gases

Group 18 elements with filled valence shells (ns2np6 except He 1s2); chemically inert under normal conditions.

Alkali Metals

Group 1 elements with ns1 valence configuration; highly reactive, form +1 ions and strong bases in water.

Alkaline-Earth Metals

Group 2 elements with ns2 valence configuration; reactive metals forming +2 ions.

Halogens

Group 17 non-metals with ns2np5 configuration; highly electronegative and form –1 ions.

Chalcogens

Group 16 elements (O, S, Se, Te, Po) characterized by ns2np4 valence shells.

Pnicogens

Group 15 elements (N, P, As, Sb, Bi) with ns2np3 outer configuration.

Shielding (Screening) Effect

Reduction of nuclear attraction on a valence electron due to repulsion by inner-shell electrons.

Effective Nuclear Charge (Z*)

Net positive charge experienced by a valence electron, calculated as Z* = Z – σ where σ is the screening constant.

Covalent Radius

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

Ionic Radius

Effective distance from the nucleus to the outermost electron of an ion; cations are smaller, anions larger, than their atoms.

van der Waals Radius

Half the distance between nuclei of adjacent, non-bonded atoms in a molecular crystal; largest operational atomic radius.

Ionisation Energy (Ionisation Potential)

Energy required to remove one electron from an isolated gaseous atom (first IE) or successive electrons (second, third IE).

Electron Affinity

Energy released when an isolated gaseous atom gains an electron to form an anion; indicates tendency to form negative ions.

Electronegativity

Tendency of an atom in a compound to attract the shared pair of electrons toward itself.

Pauling Electronegativity Scale

Most widely used quantitative scale of electronegativity based on excess bond energy differences; F assigned 4.0.

Mulliken Electronegativity

Electronegativity value obtained as the arithmetic mean of an atom’s ionisation energy and electron affinity.

Allred–Rochow Electronegativity

Scale relating electronegativity to effective nuclear charge divided by the covalent radius.

Metallic Character

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

Diagonal Relationship

Similarity of properties between certain diagonally adjacent pairs of second- and third-period elements (e.g., Li–Mg, Be–Al).

Atomic Volume

Volume occupied by one mole of atoms in the solid state; calculated as atomic weight divided by density.

Paramagnetism

Attraction of a substance to a magnetic field due to presence of unpaired electrons; magnetic moment > 0 BM.

Diamagnetism

Weak repulsion from a magnetic field exhibited by substances with all electrons paired.

Ferromagnetism

Strong magnetic ordering found in solids like Fe, Co, Ni that can retain magnetism after the external field is removed.

Bohr Magneton (BM)

Unit of magnetic moment; spin-only magnetic moment is μ = √[n(n + 2)] BM where n = number of unpaired electrons.

Hydration Energy

Enthalpy change when one mole of gaseous ions dissolves in water; increases with higher charge and smaller ionic radius.

Hydrides

Binary compounds of hydrogen; become more covalent across a period and more ionic down a group.

Oxidation Number

Effective charge an atom appears to have in a compound, representing electrons lost (positive) or gained (negative).

Amphoteric Oxide

Oxide that can react both as an acid and a base; appears near the metal-non-metal boundary (e.g., Al2O3).

Bridge Elements

Second-period elements (Li–F) whose properties resemble those of the diagonally adjacent third-period elements.

Anomalous First-Row Behavior

Distinct chemical behavior of the lightest member of a group (e.g., Li, Be) due to small size, high electronegativity and lack of d orbitals.