CHM 111: History and Development of Atomic Structure

Vocabulary practice flashcards covering the historical development of atomic models, radiation laws, quantum mechanics, and electron configurations from CHM 111.

Democritus and Leucippus

Ancient Greek philosophers who conceived the atom as an abstract thought, proposing that matter is composed of atoms without knowing their physical nature.

Dalton's Atomic Theory

A foundational atomic model proposed in 1803 based on experimentation, asserting that all chemical elements consist of tiny, indivisible, and characteristic particles called atoms.

Cathode Ray Tube Experiment

An 1896 experiment conducted by J. J. Thomson using high voltage on metals to emit identical negatively charged particles, determining their charge-to-mass ratio and discovering the electron.

Plum Pudding Model

An atomic model proposed by J. J. Thomson where negatively charged electrons are embedded uniformly within a continuous sphere of positively charged matter.

Rutherford Nuclear Atom Model

An atomic model established from alpha-particle scattering through gold foil, demonstrating that an atom possesses a tiny, dense, positively charged nucleus surrounded by electrons orbiting in mostly empty space.

Neutron

An electrically neutral subatomic particle in the atomic nucleus theorized by Ernest Rutherford and discovered experimentally by James Chadwick in 1932 to explain nuclear stability.

Continuous Spectrum

An unbroken spectrum of all wavelengths across the visible light spectrum produced when white light is dispersed through a prism.

Quantum of Energy

A discrete unit or small packet of energy emitted discontinuously by vibrating molecules of a radiating body, as proposed by Max Planck in 1900.

Planck's Constant ($h$)

A fundamental constant of proportionality equal to $6.6260755 \times 10^{-34}\,\text{J}\cdot\text{s}$ that relates the energy of a photon to its frequency via the equation $E = h\nu$.

Photoelectric Effect

The emission or ejection of electrons from the surface of a metal when irradiated with incident light of sufficient frequency, explained by Albert Einstein in 1905 using Planck's quantum hypothesis.

Threshold Frequency ($\nu_0$)

The minimum frequency of incident light required to eject an electron from a specific metal surface in the photoelectric equation $h\nu = h\nu_0 + \frac{1}{2}mv^2$.

Bohr Atomic Model

A 1913 planetary model of the hydrogen atom where electrons revolve around the nucleus only in discrete stationary circular orbits without radiating energy unless transitioning between energy levels.

Bohr Radius ($a_0$)

The radius of the lowest allowed circular orbit ($n = 1$) in the Bohr hydrogen atom, equivalent to $5.292 \times 10^{-11}\,\text{m}$ or $0.5292\,\text{\AA}$.

Rydberg Formula

An empirical and theoretical mathematical equation used to calculate the wavelengths of spectral lines of hydrogen: $\frac{1}{\lambda} = R_H \left(\frac{1}{n_1^2} - \frac{1}{n_2^2}\right)$, where $R_H = 1.097 \times 10^7\,\text{m}^{-1}$.

Lyman Series

A series of spectral emission lines of hydrogen occurring in the ultraviolet region, resulting from electron transitions dropping to the ground state level $n_1 = 1$ from higher states ($n_2 \ge 2$).

Balmer Series

A series of hydrogen spectral lines in the visible region produced by electron transitions dropping to the lower energy level $n_1 = 2$ from states $n_2 \ge 3$.

Paschen Series

A series of hydrogen emission lines in the infrared region caused by electron transitions terminating at energy level $n_1 = 3$ from levels $n_2 \ge 4$.

Moseley's Law

The relationship discovered by Henry Moseley in 1913 through characteristic X-ray spectra showing that element ordering on the periodic table is determined by atomic number rather than atomic mass.

de Broglie Hypothesis

The postulate stating that moving matter exhibits wave-like properties, with a characteristic wavelength given by $\lambda = \frac{h}{p}$, where $p$ is momentum.

Davisson-Germer Experiment

An electron diffraction experiment conducted by Clinton Davisson and Lester Germer demonstrating the wave properties and diffraction behavior of electrons.

Heisenberg Uncertainty Principle

A fundamental principle stating that it is impossible to simultaneously measure both the exact momentum and position of a subatomic particle, expressed mathematically as $\Delta p \Delta x \ge \frac{1}{2}\hbar$.

Atomic Orbital

A region in three-dimensional space around a nucleus where there is a high probability of finding an electron, mathematically described by a wave function $\psi$.

Schrödinger Wave Equation

A foundational quantum mechanical equation developed by Erwin Schrödinger relating the amplitude and spatial wave function $\psi$ of an electron to its total and potential energy.

Dirac Relativistic Quantum Mechanics

The 1928 theoretical reformulation of quantum mechanics by Paul A. M. Dirac incorporating special relativity, leading directly to the introduction of the electron spin quantum number.

Principal Quantum Number ($n$)

The primary quantum number designating the main shell or energy level of an electron ($n = 1, 2, 3, \dots$), dictating a maximum capacity of $2n^2$ electrons per shell.

Angular Momentum Quantum Number ($l$)

The quantum number designating orbital shape and subshell type ($s, p, d, f$), taking integral values from $0$ up to $(n - 1)$.

Magnetic Quantum Number ($m_l$)

The quantum number designating the specific orientation of an orbital in three-dimensional space, with integral values ranging from $-l$ through $0$ to $+l$.

Spin Quantum Number ($m_s$)

The quantum number representing the intrinsic angular momentum and magnetic field orientation of an electron, taking values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

Aufbau Principle

The rule stating that electrons fill available atomic orbitals in order of increasing energy levels, generally filling subshells with lower $(n + l)$ values first.

Pauli Exclusion Principle

The quantum mechanical rule stating that no two electrons in the same atom can possess identical sets of all four quantum numbers ($n, l, m_l, m_s$).

Hund's Rule

The rule dictating that degenerate orbitals in a subshell are each occupied singly with parallel electron spins before any orbital is doubly occupied.