Electronic Structure of Atoms

Vocabulary flashcards covering key terms and concepts related to the electronic structure of atoms.

Electromagnetic Radiation (EMR)

Electromagnetic waves have characteristic wavelengths () and frequencies (). Understanding EMR is essential to understand the electronic structure of atoms.

Wavelength (λ)

The distance between corresponding points on adjacent waves.

Frequency (ν)

The number of cycles that pass a point in one second; measured in Hertz (Hz).

Speed of Light (c)

All types of electromagnetic radiation (EMR) move through a vacuum at the speed of light c, which is approximately 3.00 × 10^8 m/s.

Relationship between Wavelength and Frequency

Wavelength () and frequency () have an inverse relationship: as wavelength increases, frequency decreases, and vice versa. Expressed as c = 

Planck's Constant (h)

A fundamental constant that relates the energy of a photon to its frequency; approximately 6.626 × 10–34 J.s.

Energy of a Photon (E)

The energy of a photon is directly proportional to its frequency and inversely proportional to its wavelength. E = hν = hc/λ

Electromagnetic Spectrum

The range of all types of electromagnetic radiation; high energy radiations (X-rays) have shorter wavelengths than low energy radiations (radio waves).

Quantum

The smallest amount of energy (photon) that can be emitted or absorbed as electromagnetic radiation.

Photoelectric Effect

The emission of electrons from metal surfaces on which light shines; provides evidence for the particle nature of light and quantization.

Photons

Energy packets or quanta of light.

Quantized Energy

Energy exists at discrete levels; electrons must gain a specific amount of energy to move from one energy level to another.

Bohr's Model

Model of the atom where electrons are confined to specific energy states or orbits.

Principal Quantum Number (n)

Describes the energy level or shell on which the orbital resides; integer values ≥ 1.

Ground State

The lowest energy state of an atom (n=1).

Absorption (Excitation)

The process where an electron gains energy and moves to a higher energy level.

Emission

The process where an electron loses energy and moves to a lower energy level, releasing a photon.

Quantum Mechanics

Mathematical treatment that incorporates both the wave and particle nature of matter.

Orbitals

Spatial distribution of electron density described by a set of three quantum numbers (n, l, ml).

Azimuthal (Angular) Quantum Number (l)

Defines the shape of the orbital; values range from 0 to (n − 1). l = 0, 1, 2, 3 corresponds to s, p, d, f orbitals, respectively.

Magnetic Quantum Number (ml)

Describes the three-dimensional orientation of the orbital in space; values are integers ranging from -l to +l.

Degenerate Orbitals

Orbitals on the same energy level that have the same energy (only in one-electron systems like hydrogen).

Spin Quantum Number (ms)

Describes the intrinsic angular momentum of an electron, which is quantized and has two possible values: +1/2 (spin up) and −1/2 (spin down).

Pauli Exclusion Principle

No two electrons in the same atom can have identical sets of four quantum numbers (n, l, ml, ms).

Electron Configuration

Distribution of all electrons in an atom among the various orbitals and energy levels.

Hund's Rule

For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.

Valence Electrons

Electrons in the outermost shell of an atom; these electrons are involved in chemical reactions.

Condensed Electron Configurations

Shorthand notation for electron configurations where core electrons are represented by the symbol of the preceding noble gas in square brackets.

Isoelectronic

Describes atoms or ions that have identical electron configurations.