Atomic Theory, Atomic Structure, and the Periodic Table — Vocabulary Flashcards

A comprehensive set of vocabulary flashcards covering atoms, atomic structure, periodic table concepts, and fundamental models of the atom from the provided notes.

Atom

The smallest part of an element that has all the properties of that element; the basic unit of matter.

Mass

Quantity of matter in a body; independent of volume and forces acting on it.

Weight

The force of gravity acting on a body.

Nucleus

The dense, positively charged center of the atom that contains protons and neutrons.

Electron

Tiny, negatively charged particle that orbits the nucleus; first discovered by Thomson (1897).

Proton

Positively charged particle in the nucleus; determines the element; mass about 1 amu.

Neutron

Neutral particle in the nucleus; heavier than an electron; no electric charge.

Dalton (1808)

Developed modern atomic theory: atoms are indivisible, atoms of the same element are identical, and atoms combine in simple whole-number ratios.

J.J. Thomson

Proposed the Plum Pudding model; introduced electrons as corpuscles and showed atoms contain smaller charged parts.

Plum Pudding Model

Atom described as a positive pudding with negatively charged electrons embedded (plums).

Rutherford (1911)

Proposed the nuclear model: a tiny, dense, positively charged nucleus with electrons around it.

Nucleus (Rutherford)

Center of the atom containing protons and neutrons; mass concentrated there.

Bohr (1913)

Proposed the planetary model: electrons occupy specific energy levels (orbits) around the nucleus.

Planetary Model

Electrons move in fixed orbits around the nucleus at certain energy levels.

Schrödinger (1926)

Developed the quantum mechanical model: electrons described by probability densities (electron cloud).

Wave Model

Electrons do not have exact paths; location is probabilistic and energy-dependent; described by wave mechanics.

Electron Configuration

Arrangement of electrons in shells, subshells, and orbitals according to energy levels.

Shell

Energy levels around the nucleus where electrons reside; closer shells are lower in energy.

Subshells

Divisions within a shell, denoted by s, p, d, f.

Orbitals

Probability regions around the nucleus that can hold up to 2 electrons with opposite spins.

Pauli Exclusion Principle

No two electrons in the same orbital can have identical quantum numbers; two electrons per orbital with opposite spins.

Valence Electron

Electrons in the outermost shell that determine an atom’s bonding and reactivity.

Electron Binding Energy (Eb)

Strength with which an electron is held to the nucleus; higher for inner electrons; measured in eV.

Nuclide

An atomic species with a definite number of protons and neutrons.

Isotopes

Nuclides of the same element with the same Z but different mass number A.

Isobars

Nuclides with the same mass number A but different atomic numbers Z.

Isotones

Nuclides with the same neutron number N but different Z.

Isomers

Nuclides with the same Z and A but different energy states and spins.

Metastable State

Isomers with half-lives long enough to be considered in a stable state (typically ≥ 10^-9 s).

Atomic Number (Z)

Number of protons in the nucleus; defines the element and its position in the periodic table.

Atomic Mass (A)

Mass of an atom in atomic mass units (amu); roughly the sum of protons and neutrons.

Periodic Table

Tabular display of elements arranged by atomic number, electron configuration, and recurring properties; introduced by Mendeleev.

Dmitri Mendeleev

Chemist who formulated the periodic table in 1869 and predicted unknown elements.

Groups (Families) in PTE

Vertical columns in the periodic table; elements in a group share similar properties.

Periods in Periodic Table

Horizontal rows; indicate energy levels of valence electrons (seven periods).

Blocks in Periodic Table

Regions (S, P, D, F) defined by the orbital type being filled.

S-Block

Left side of the table (IA and IIA, plus H and He); elements with s-orbital electrons being filled.

P-Block

Right side; groups 3A-8A; contains metals, nonmetals, and metalloids.

D-Block

Transition metals; elements between s- and p-blocks.

F-Block

Inner transition metals (lanthanides and actinides); f-block elements.

Electron Configuration – Shells

Distribution of electrons among shells with increasing energy away from the nucleus.

Electron Configuration – Subshells

Within each shell, electrons occupy subshells denoted by s, p, d, f.

Electron Configuration – Orbitals

Specific regions within subshells that hold up to 2 electrons with opposite spins.

Shell vs Subshell vs Orbital

Shell = energy level; Subshell = s/p/d/f subdivision; Orbital = max 2 electrons in a region.

Valence Electrons and Bonding

Electrons in the outermost shell that determine bonding; full outer shell means low reactivity.

Atomic Nomenclature – Nuclide/Isotopes/Isobars/Isotones/Isomers

Terminology to describe nuclear species: nuclide, isotopes, isobars, isotones, and isomers.

Alkali Metals (Group 1)

Very reactive metals; react vigorously with water.

Alkali Earth Metals (Group 2)

Reactive metals with two valence electrons; lose both to achieve stability.

Lanthanides

Rare earth metals; silvery; occur sparsely in Earth's crust.

Actinides

Radioactive metals; thorium and uranium occur naturally; others are synthetic.

Transition Metals

Hard, shiny, malleable metals; good conductors; include iron, copper, silver, gold.

Post-Transition Metals

Elements between transition metals and metalloids; softer and poorer conductors.

Metalloids

Elements with properties between metals and nonmetals; semi/poor conductors.

Non-Metals

Mostly gases; include hydrogen; generally poor conductors.

Halogens (Group 17)

Highly reactive nonmetals; form salts with alkali metals.

Noble Gases (Group 18)

Colorless, odorless, inert/non-reactive gases.