Classification of Elements and Periodicity in Properties

Vocabulary flashcards covering key terms and definitions from the lecture notes on the Periodic Table and periodicity.

Periodic Table

An organized arrangement of the chemical elements into groups (families) and periods that reveals trends and relationships among elements, showing that elements lie in families rather than as a random cluster.

Periodic Law (Modern)

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

Mendeleev's Periodic Law

The properties of the elements are a periodic function of their atomic weights; he left gaps for undiscovered elements and grouped elements by similar properties.

Atomic number

Z; the number of protons in the nucleus (and, for a neutral atom, the number of electrons); the fundamental parameter for periodic classification.

Electronic configuration

The distribution of electrons among atomic orbitals (s, p, d, f); the last orbital filled helps determine an element’s position in the table.

Period

A horizontal row in the periodic table; reflects the principal quantum number of the outer shell; lengths are 2, 8, 8, 18, 18, 32, with the seventh period incomplete.

Group (family)

A vertical column in the periodic table; elements in a group share similar outer electronic configurations and properties; groups are numbered 1–18 by IUPAC.

s-block elements

Elements with outermost electrons in s orbitals (ns1 or ns2); includes Group 1 (alkali metals) and Group 2 (alkaline earth metals); generally highly reactive.

p-block elements

Elements with outermost electrons in p orbitals (ns2np1 to ns2np6); Groups 13–18; includes metals, metalloids, and non-metals; ends with noble gases.

d-block elements

Transition elements; Groups 3–12; characterized by filling of inner d orbitals; typically metals with variable oxidation states and catalysts; many are colored ions.

f-block elements

Inner-transition elements; Lanthanoids (Ce–Lu) and Actinoids (Th–Lr); outer configuration (n−2)f1−14; typically radioactive and placed in bottom panels.

Lanthanoids

The 4f-block elements Ce–Lu; part of the inner-transition series; show similar chemistry within the series.

Actinoids

The 5f-block elements Th–Lr; inner-transition elements; many are radioactive and some have complex chemistry.

Long form periodic table

The modern arrangement with 7 periods and 18 groups; bottom panels house the lanthanoids and actinoids; groups are numbered 1–18.

Nomenclature of elements with atomic number above 100 (IUPAC)

Temporary names for new elements using three-letter roots (e.g., unnilunium) and the suffix -ium until official names are adopted.

IUPAC official names for Z>100

Permanent names and symbols assigned by IUPAC after discovery is confirmed; all elements up to Z=118 have official names.

Eka-aluminium

Mendeleev’s predicted element (gallium) placed below aluminium; demonstrated the predictive power of his table.

Eka-silicon

Mendeleev’s predicted element (germanium) placed below silicon; later discovered and matched predictions.

Dobereiner's Triads

Early grouping of elements in threes where the middle element’s properties lay between the other two; led to ideas about periodicity.

Law of Octaves

Newlands’ rule that every eighth element has similar properties to the first; valid only up to calcium and not universally accepted.

De Chancourtois

French geologist who proposed a cylindrical (tubular) arrangement of elements by atomic weights to reveal periodicity.

Moseley

Physicist who showed atomic number, not atomic weight, governs periodicity by using X-ray spectra; refined the Periodic Law.

Mendeleev

Russian chemist who published the first broadly used periodic table and predicted gaps for undiscovered elements (e.g., Eka-Al, Eka-Si).

Hydrogen

A unique element with one 1s electron; placed at the top of the table and can be treated as behaving like either Group 1 or Group 17 in some contexts; often shown separately.

Noble gases

Group 18 elements with complete valence shells; very low reactivity due to stability of closed electron configurations.

Halogens

Group 17 elements; highly reactive non-metals with high electron affinity; tend to gain one electron to achieve noble-gas configuration.

Metals

Elements on the left side of the periodic table; typically good conductors, malleable, and metallic character increases down a group.

Non-metals

Elements on the right side of the periodic table; generally poor conductors, with metallic character decreasing across a period.

Metalloids (semi-metals)

Elements with properties intermediate between metals and non-metals; lie along the diagonal boundary line in the table.

Electronegativity

Qualitative measure of an atom’s ability to attract electrons in a bond; on the Pauling scale, increases across a period and decreases down a group.

Ionization enthalpy

Energy required to remove an electron from a gaseous atom; first ionization enthalpy generally increases across a period and decreases down a group.

Electron gain enthalpy

Energy change when an electron is added to a neutral atom to form an anion; generally more negative across a period and less negative down a group; anomalies exist (e.g., O, F).