Chemistry of the Elements in the Periodic Table - VOCAB Flashcards

Vocabulary flashcards covering key concepts from the lecture on the periodic table, trends, electron configuration, and basic reactions.

Periodic Table

A tabular arrangement of elements ordered by increasing atomic number, showing periods and groups and recurring chemical properties.

Period

A horizontal row in the periodic table; represents the principal energy level for valence electrons and shows systematic changes across the table.

Group/Family

A vertical column in the periodic table; elements in a group have similar chemical properties; for main-group elements, the group number often indicates valence electrons.

Representative elements (A)

Elements in the s- and p-block (main-group elements) with properties that vary predictably with group number.

Transition metals (B)

The d-block elements with typically variable oxidation states and many colored compounds; good conductors of electricity.

Inner transition metals (f-block)

The f-block elements (lanthanides and actinides) characterized by similar properties within each series.

Lanthanide series

The 15 f-block elements from La to Lu; generally non-radioactive except Promethium.

Actinide series

The 15 f-block elements from Ac to Lr; all radioactive.

Atomic number

The number of protons in an atom's nucleus; defines the element's identity and order on the periodic table.

Atomic mass

Weighted average mass of an element's isotopes; historically used for tabulation; not simply an integer.

Law of Octaves

Early idea that properties repeat every eight elements when arranged by increasing atomic weight; not universally valid.

Octet rule

Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (stable noble-gas configuration).

Diagonal relationship

Similar properties between certain pairs of elements in adjacent periods (e.g., Li–Mg, Be–Al, B–Si).

Ionization energy

Energy required to remove the outermost electron from a neutral atom; generally increases across a period and decreases down a group.

Electron affinity

Energy released when a neutral atom gains an electron; trends vary but become more negative across a period.

Electronegativity

An atom's ability to attract electrons in a chemical bond; increases across a period and decreases down a group.

Atomic radius

Half the distance between nuclei in a diatomic molecule; increases down a group and decreases across a period.

Metallic character

Tendency of an element to exhibit metallic properties; decreases across a period and increases down a group.

Principal quantum number (n)

Main energy level of an electron; determines the size of the orbital.

Azimuthal quantum number (l)

Subsidiary quantum number describing orbital shape: s (l=0), p (l=1), d (l=2), f (l=3).

Magnetic quantum number (m_l)

Orbital orientation within a subshell; values range from -l to +l.

Spin quantum number (m_s)

Electron spin; values are +1/2 or -1/2.

Aufbau principle

Electrons fill lowest-energy orbitals first before occupying higher-energy ones.

Electron configuration

Arrangement of electrons in atoms (e.g., Ca: [Ar] 4s2).

Pauli exclusion principle

No two electrons in an atom can have the same set of quantum numbers.

Heisenberg uncertainty principle

Cannot simultaneously know exact position and momentum of a particle.

Hund's rule

Electrons in equal-energy orbitals occupy separate orbitals with parallel spins before pairing.

Noble gases

Group 8A; elements with full valence shells; chemically inert under normal conditions.

Halogens

Group 7A; seven valence electrons; highly reactive nonmetals; form -1 anions.

Alkali metals

Group 1A; one valence electron; form +1 ions; highly reactive.

Alkaline earth metals

Group 2A; two valence electrons; form +2 ions; reactive but less than group 1A.

Carbon family

Group 4A; four valence electrons; includes C, Si, Ge, Sn, Pb; multiple oxidation states.

Nitrogen family

Group 5A; five valence electrons; includes N, P, As, Sb, Bi; multiple oxidation states.

Oxygen family

Group 6A; six valence electrons; includes O, S, Se, Te, Po; commonly -2 oxidation state.

Transuranic elements

Elements with atomic number greater than 92; radioactive; produced synthetically.

Nihonium (Nh)

Element 113.

Moscovium (Mc)

Element 115.

Tennessine (Ts)

Element 117.

Oganesson (Og)

Element 118.

Glenn Seaborg

Chemist who discovered transuranic elements and helped shape the actinide concept.

Haber’s process

Synthesis of ammonia from nitrogen and hydrogen: 3H2 + N2 → 2NH3.

Synthesis (combination) reaction

A + B → AB; two substances combine to form a compound.

Decomposition (analysis) reaction

AB → A + B; a compound breaks down into simpler substances.

Single displacement

AB + X → AX + B; more reactive element displaces another from a compound.

Double displacement

AB + CD → AD + CB; exchange of partners between two compounds.

Neutralization

Acid-base reaction producing salt and water, e.g., NaOH + HCl → NaCl + H2O.

Precipitation

Formation of an insoluble solid (precipitate) when solutions react.

Combustion

Fuel reacts with O2 to form CO2 and H2O (complete combustion).

Electrolysis

Electrical current drives decomposition of a compound, e.g., 2H2O → 2H2 + O2.

Periodic law

Properties of elements are periodic functions of their atomic number.