Electromagnetic Radiation and Atomic Models

This set of vocabulary flashcards covers the electromagnetic spectrum, wave properties, light energy calculations, and the transition from the Bohr model to the Quantum Mechanical model of the atom, including electron configurations and periodic trends.

Electromagnetic radiation

A form of energy that travels in waves and exists in increments called photons.

Wavelength ($\lambda$)

The length of one wave, often measured in meters ($m$) or nanometers ($nm$).

Frequency ($\nu$)

The number of waves per second, where $1\,\text{wave/second} = 1\,\text{hertz (Hz)}$.

Speed of Light ($c$)

The product of wavelength and frequency, equal to a constant value of $3.00 \times 10^8\,\text{m/s}$.

Visible Spectrum

A small range of electromagnetic radiation that human eyes can detect, roughly spanning from $400\,\text{nm}$ to $700\,\text{nm}$.

Planck’s constant ($h$)

A physical constant used to calculate the energy of a photon, equal to $6.63 \times 10^{-34}\,J \cdot s$.

Energy of a photon ($E$)

Calculated using the equations $E = h\nu$ or $E = \frac{hc}{\lambda}$, showing energy is directly proportional to frequency and inversely proportional to wavelength.

Line Spectra

Distinct lines of colors produced when light from a gas lamp passes through a prism, used as "fingerprints" to identify elements.

Bohr Model

A 1913 model proposing that electrons orbit the nucleus in specific "allowed" energy levels and can jump between levels by absorbing or releasing light.

Ground state

The state of an atom where all electrons are in the lowest possible energy levels.

Heisenberg’s Uncertainty Principle

The principle stating that it is impossible to precisely know the exact velocity and location of a particle simultaneously.

Quantum Model

A model developed in the 1920s-30s that describes electrons as having wave-like nature and occupying regions of space called orbitals based on probability.

Principal quantum number ($n$)

An integer ($1, 2, 3...$) used to identify the main energy levels occupied by electrons.

Sublevel

Divisions within energy levels labeled $s, p, d,$ and $f$, each with a different electron capacity and number of orbitals.

Orbital

The region around the atom where an electron is most likely to be found; each orbital can hold up to two electrons.

Spin

A property of electrons that creates a magnetic field; electrons with opposite spins pair together within an orbital.

Hund’s Rule

The rule stating that if empty orbitals of the same energy are available, electrons will singly occupy them before pairing up.

Valence level

The highest occupied electron energy level in an atom, which can hold up to 8 electrons.

Octet Rule

The principle that an atom is stabilized by having its highest-occupied (valence) energy level filled.

Noble gas shorthand

A method of writing electron configurations where the symbol of the previous noble gas in brackets (e.g., $[Ne]$) represents the inner-shell electrons.

Isoelectronic

A term describing different atoms or ions that possess the same electron configuration.

Periodic Table Blocks

The categorization of elements into $s$, $p$, $d$, and $f$ blocks based on the highest energy sublevel their electrons occupy.