Atomic Theory and Periodic Trends

Vocabulary flashcards covering the history of atomic theory, quantum mechanics, electron configurations, and periodic trends in properties and chemical bonding.

Democritus

A Greek philosopher (430 BCE) who reasoned that all matter is composed of tiny, indestructible particles in empty space called "atomos" (uncuttable).

Dalton's Atomic Theory

An 1803 model proposing that all matter consists of indivisible atoms, atoms of a given element are identical, atoms combine in whole-number ratios to form compounds, and chemical reactions separate, combine, or rearrange atoms.

Cathode Ray Tube

A sealed vacuum glass tube developed by William Crookes that emits beams of negatively charged particles, which led to the discovery of the electron.

Electron

The first subatomic particle discovered (by J.J. Thomson in 1897), consisting of a negatively charged particle with a mass significantly smaller than that of an atom.

Plum-Pudding Model

An atomic model proposed by J.J. Thomson in 1903 featuring a positively charged sphere embedded throughout with negatively charged electrons.

Gold Foil Experiment

A 1909 experiment by Ernest Rutherford where alpha particles were directed at thin gold foil; the deflection of a small fraction proved the existence of a tiny, dense, positively charged nucleus.

Nucleus

The tiny, dense, positively charged core at the center of an atom discovered by Ernest Rutherford.

Neutron

A neutral subatomic particle located in the nucleus that contributes mass without charge, proved to exist by James Chadwick in 1932.

Photoelectric Effect

The phenomenon explained by Albert Einstein in 1905 whereby energy packets called photons eject electrons from a material when the radiation has sufficient energy or frequency.

Bohr Model

A planetary model of the atom proposed in 1913 by Niels Bohr, stating that electrons occupy distinct, stable energy levels around the nucleus without radiating energy.

Line Spectrum

A series of discrete wavelengths or colors of light emitted by the excited electrons of a specific gas when returning to their ground state.

Valence Shell

The outermost occupied electron shell of an atom containing the electrons responsible for chemical reactivity.

Valence Electrons

Electrons residing in the outermost occupied shell of an atom that participate in bonding.

Heisenberg Uncertainty Principle

A quantum mechanical principle formulated in 1925 stating that it is impossible to determine simultaneously both the exact position and momentum of a subatomic particle.

Atomic Orbital

A three-dimensional region of space around the nucleus mapped by Erwin Schrödinger where there is a high probability of finding an electron.

Aufbau Principle

The rule stating that electrons fill available orbitals of lowest energy first, progressing through sublevels in order of increasing energy ($s < p < d < f$).

Hund's Rule

The principle stating that electrons occupy degenerate orbitals singly with parallel spins before pairing, minimizing electron-electron repulsion.

Pauli Exclusion Principle

The quantum rule dictating that no two electrons in an atom can have the exact same four quantum numbers, requiring paired electrons in the same orbital to have opposite spins.

Core Notation

A condensed electron configuration shorthand in which the configuration of the preceding noble gas is symbolized in square brackets, followed by the remaining valence electrons.

Internuclear Distance

The distance between the centers of two bonded atomic nuclei, used to determine atomic and ionic radii.

Atomic Radius

Defined as half the internuclear distance between two identical bonded atoms ($\text{Atomic Radius} = \frac{\text{internuclear distance}}{2}$); it increases down a group and decreases left-to-right across a period.

Effective Nuclear Charge ($Z_{eff}$)

The net positive charge experienced by valence electrons from the nucleus; an increase in $Z_{eff}$ draws outer electrons closer, decreasing atomic size across a period.

Shielding Effect

The phenomenon whereby inner electron shells obstruct the nuclear attraction felt by outer valence electrons, allowing atoms to increase in size down a group.

Ionization Energy

The minimum energy required to remove an electron from a gaseous atom or ion in its ground state.

Electronegativity

The relative ability of a bonded atom to attract shared electrons toward its own nucleus.

Non-Polar Covalent Bond

A chemical bond characterized by equal sharing of electrons between bonded atoms, occurring when the electronegativity difference is negligible ($\Delta EN = 0$).

Polar Covalent Bond

A bond formed when electrons are shared unequally between atoms with an electronegativity difference between $0.4$ and $1.7$, creating partial positive ($\delta+$) and partial negative ($\delta-$) charges.

Ionic Bond

A bond characterized by the complete transfer of one or more valence electrons between atoms with an electronegativity difference greater than $1.7$ ($\Delta EN > 1.7$), creating cations and anions.

Lewis Dot Diagram

A two-dimensional representation of an atom or ion using its chemical symbol flanked by dots that represent only its valence electrons.

Electron-Deficient Molecule

A stable molecule containing central atoms such as $H$, $He$, $Be$, $B$, or $Al$ that possess fewer than eight valence electrons, representing an exception to the octet rule.

Expanded Octet

An exception to the octet rule occurring in nonmetals of period 3 or higher, where central atoms accommodate more than eight valence electrons by utilizing unoccupied $d$-orbitals.