Electronic Structure of Atoms

This set of vocabulary flashcards covers the historical and theoretical developments of the electronic structure of atoms, including wave-particle duality, quantum mechanics, orbital filling rules, and periodic properties.

Wavelength ($\lambda$)

The distance between identical points on a wave, measured in meters.

Frequency ($\nu$)

The number of waves that pass through a certain point in 1 second, measured in Hz ($1\,\text{cycle/s}$).

William Herschel

The scientist who discovered "invisible" or Infrared (IR) light in 1800.

Double Slit Interference Experiment

An 1801 experiment by Thomas Young that proved light travels as a wave due to the observation of multiple wave interactions known as interference.

James Maxwell

The 1873 scientist who determined that light consists of interconnected electricity and magnetism traveling via electromagnetic waves.

Speed of light ($c$)

The constant speed at which all electromagnetic light radiation travels, equal to $3.0 \times 10^8\,m/s$ or approximately $670\,\text{million mph}$.

Max Planck

The scientist who solved the "Heated Solids Problem" in 1900 by proposing that energy is quantized and must be transferred in precise "quantities" ($\Delta E = h \times \nu$).

Planck's constant ($h$)

A physical constant used to describe the sizes of quanta, equal to $6.63 \times 10^{-34}\,J \cdot s$.

Photoelectric Effect

The phenomenon solved by Einstein in 1905 where electrons are ejected from metal only when exposed to a certain minimum frequency of light, proving light's particle nature.

Photon

A massless "particle" or quantized stream of light proposed by Albert Einstein.

Spectroscopy

The study of the interaction between light and matter.

Atomic line spectrum

A spectrum produced by elements when heated that shows only specific changes in energy ($\Delta E$) as distinct lines of visible radiation, rather than a continuous spectrum.

Niels Bohr

The scientist who determine in 1913 that atomic line spectra occur due to distinct electron transitions between specific energy shells.

Ground state

The lowest possible energy level for an electron ($n = 1$).

Excited state

A state in which an electron has absorbed energy and been promoted to a shell higher than its ground state.

Flame test

A qualitative test used to identify metal ions in compounds based on the unique wavelengths of light released as electrons relax to the ground state.

Rydberg constant ($R_H$)

A constant used in the calculation of the energies of a Hydrogen electron, equal to $2.18 \times 10^{-18}\,J$.

L.A.S.E.R.

An acronym for Light Amplification by Stimulated Emission of Radiation, producing intense, mono-energetic (single wavelength), and coherent light.

de Broglie’s hypothesis

The 1923 proposal by Louise de Broglie that electrons, like light, have both particle and wave properties (wave-particle duality).

Heisenberg Uncertainty Principle

The 1927 principle stating it is impossible to know both the momentum ($\text{mass} \times \text{velocity}$) and the position of a particle like an electron with absolute certainty.

Quantum Model of the Atom

A model that describes electron locations as probability distributions, often referred to as an "electron cloud."

Wave function ($\psi$)

A set of equations from Erwin Schrodinger that describes the energy of an electron and the probability of finding it in a specific volume of space.

Orbital

An area of high probability (where 90% of the electron density is found) for locating an electron.

Principal quantum number ($n$)

The quantum number that describes the main energy level or shell ($n = 1, 2, 3, \dots$) and the distance of the electron from the nucleus.

Angular momentum quantum number ($l$)

The quantum number that defines the shape of the electron orbital (s, p, d, or f) where $l = 0, \dots, n-1$.

Magnetic Quantum Number ($m_l$)

The quantum number that distinguishes separate orbitals within a shell of the same type ($l$), with values ranging from $-l$ to $+l$.

Spin quantum number ($m_s$)

The quantum number describing the quantum spin of an electron, which can be either $+1/2$ or $-1/2$.

Pauli exclusion principle

A principle stating that no two electrons in an atom can have the same four quantum numbers ($n, l, m_l, m_s$).

Aufbau principle

The rule stating that atoms "fill up" the lowest energy electron orbitals first before moving to higher energy subshells.

Hund’s Rule

The rule stating that the most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins; electrons will not double up in an orbital until necessary.

Paramagnetic

A state where magnetic fields are not cancelled because electrons are not paired; these materials are potentially magnetic.

Diamagnetic

A state where opposite magnetic fields cancel each other out because all electrons are paired.

Valence electrons

The outer shell electrons of an atom located in the highest energy shell ($n$) that participate in chemical bonding.

Octet Rule

The tendency of atoms to form bonds or ionize to achieve a full valence shell of 8 electrons, similar to a Noble Gas.

Isoelectronic

A term describing species that have the same number of electrons and the same electron configuration (e.g., $Ne$, $N^{-3}$, and $Mg^{+2}$).