Periodic Table, Periodic Properties and Variations of Properties (Vocabulary Flashcards)

A set of vocabulary-style flashcards covering the key terms and concepts from the lecture notes on periodic properties and the modern periodic table.

Periodic law

The properties of elements are periodic functions of their atomic numbers (as discovered by Moseley); elements show repeating patterns when arranged by increasing atomic number.

Modern periodic table

The long-form table with 18 groups and 7 periods, organized by increasing atomic number; groups are numbered 1–18 and elements are placed in s-, p-, d-, or f-blocks.

Group

A vertical column in the periodic table; elements in a group have similar valence electron configurations and similar chemical properties.

Period

A horizontal row in the periodic table; the number of shells (electronic levels) increases as you move down a period.

Representative elements

Elements in groups 1, 2, and 13–18 (s and p blocks) known for typical main-group properties.

Transition elements

Elements in groups 3–12 (d-block) with incomplete d subshells; exhibit characteristic transition metal properties.

Noble gases

Group 18 elements with a complete outer electron shell and very low chemical reactivity.

Alkali metals

Group 1 elements; highly reactive metals that lose one electron to form positive ions and react vigorously with water.

Alkaline earth metals

Group 2 elements; form weaker bases than alkali metals; lose two electrons to form cations.

Halogens

Group 17 elements; highly reactive non-metals; gain one electron to complete octet and form salts.

Lanthanide series

Inner transition elements in period 6 (f-block, 57–71); similar properties due to filling of 4f subshell.

Actinide series

Inner transition elements in period 7 (f-block, 89–103); many are radioactive; 5f subshell involvement.

Inner transition elements

The f-block elements (lanthanides and actinides) sitting below the main body of the periodic table.

Atomic radius (atomic size)

Half the distance between the nuclei in a diatomic molecule; distance from the nucleus to the outermost shell.

Nuclear charge

The positive charge of the nucleus equal to the number of protons (the atomic number); influences attraction on electrons.

Ionisation energy (ionisation potential)

Energy required to remove one electron from a neutral gaseous atom to form a cation (IE1; units: kJ/mol or eV/atom).

Electron affinity

Energy change when adding an electron to a neutral gaseous atom to form an anion; often negative for many elements, with halogens having large negative values.

Electronegativity

Tendency of an atom to attract a shared electron pair in a chemical bond; Pauling scale commonly used (F = 4.0, Cs ≈ 0.7).

Valence electrons

Electrons in the outermost shell that participate in bonding and determine chemical properties.

Valency

The combining capacity of an atom; often the number of electrons it can donate, accept, or share to attain a stable octet.

Octet rule

Atoms tend to gain, lose, or share electrons to achieve a full outer shell of eight electrons.

Isoelectronic ions

Different species that have the same number of electrons; their sizes depend on nuclear charge.

Diagonal relationship

Similarities between certain elements in adjacent periods (e.g., Li–Mg, Be–Al, B–Si) due to small electronegativity differences.

Periodicity

Regular reappearance of properties at intervals because of the recurring electronic configurations in outer shells.

Shells (K, L, M, N, O, P, Q)

The named electron shells corresponding to n = 1 to n = 7; K = 1st shell, L = 2nd, etc.

Oxides: basic, amphoteric, acidic

Trends show basic oxides for metals, amphoteric oxides for some metals/metalloids, and acidic oxides for non-metals, changing across periods.

Hydrogen as a special nonmetal

Hydrogen is a nonmetal with one electron and one proton; can form H+ and H- under certain conditions.

Electronic configuration

Arrangement of electrons in shells around the nucleus (e.g., Na = 2, 8, 1; Cl = 2, 8, 7).

Long form periodic table

The version of the periodic table showing the complete layout, including Lanthanides and Actinides at the bottom.

IUPAC group numbering

Groups are numbered from 1 to 18 in modern practice (replacing the old IA–VIIA, VIII, etc.).

Old group notation

Earlier labeling (IA, IIA, IIIB, etc.) used in the study material; often accepted with the new numbering.

Noble gas electron configuration impact

Noble gases have complete octets and do not tend to gain electrons; zero electron affinity for stable configurations.

Atomic number (Z)

Number of protons in the nucleus; unique identifier for an element and equals the number of electrons in a neutral atom.

Mass number (A)

Total number of protons and neutrons in the nucleus; A = Z + number of neutrons.

Isotopes

Atoms of the same element with the same Z but different A due to varying neutron numbers; differ in mass.

Electronic configuration pattern across periods

Across a period, adding protons increases nuclear pull, reduces atomic size, and increases electronegativity and ionisation energy.

Metallic character

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

Non-metallic character

Tendency to gain electrons; generally decreases down a group and increases across a period.

Dielectric trend in reactivity

Metals: reactivity increases down a group; non-metals: reactivity generally decreases down a group.

Electronegativity trend across periods

Electronegativity increases across a period due to increasing nuclear charge and decreasing atomic size.