CEM 141 - Chapter 2: Electrons and Orbitals

Vocabulary practice flashcards covering wave-particle duality, atomic spectroscopy, quantum mechanical atomic models, electron configurations, and periodic trends from Chemistry Chapter 2.

Wavelength ($\lambda$)

The distance between two consecutive peaks or troughs in a wave, measured in meters ($\text{m}$) or nanometers ($\text{nm}$).

Frequency ($\nu$)

The number of complete wave cycles that pass a given point per unit time, measured in Hertz ($\text{Hz} = \text{s}^{-1}$).

Speed of Light ($c$)

The constant speed at which all electromagnetic radiation travels in a vacuum, equal to $3.00 \times 10^8\,\text{m/s}$, relating wavelength and frequency through $c = \lambda \nu$.

Amplitude

The height of a wave from its baseline to a peak (or trough); in the wave model of light, it corresponds to intensity/brightness, but it does not determine the energy of individual photons.

Diffraction

The bending and spreading of waves as they encounter an obstacle or pass through an aperture/slit; particles do not diffract, making diffraction definitive evidence of wave-like behavior.

Constructive Interference

The interaction that occurs when two waves combine in-phase such that their crests and troughs reinforce one another, producing an increased overall amplitude.

Destructive Interference

The interaction that occurs when two waves combine out-of-phase such that the crest of one wave cancels the trough of another, yielding diminished or zero amplitude.

Photoelectric Effect

The ejection of electrons from a metal surface when light shines upon it, demonstrating that light transfers discrete packets of energy directly to electrons.

Threshold Frequency

The minimum frequency of incident radiation required to eject electrons from a specific metal surface during the photoelectric effect, below which no electrons are emitted regardless of light intensity.

Photon

A discrete quantum packet or particle of electromagnetic energy, with energy defined by the relation $E = h\nu$.

Planck's Constant ($h$)

The proportionality constant relating the energy of a photon to its frequency, with an experimental value of $6.626 \times 10^{-34}\,\text{J}\cdot\text{s}$.

Continuous Spectrum

A radiation spectrum containing all uninterrupted wavelengths and colors over a given region, produced by white light sources like the Sun.

Atomic Emission Spectrum

A series of discrete, specific bright lines of light emitted by excited atoms of an element as electrons transition from higher to lower energy levels.

Atomic Absorption Spectrum

A series of discrete dark lines superimposed on a continuous spectrum, produced when cold gaseous atoms absorb specific wavelengths of light corresponding to electronic transitions from lower to higher energy levels.

Bohr Model

An atomic model postulating that electrons orbit the nucleus in specific circular paths with discrete, quantized energies and radii; it accurately describes the hydrogen atom but fails for multi-electron systems.

Ionization

The process wherein an atom absorbs sufficient energy to completely remove an electron, effectively promoting the electron from its bound state to energy level $n = \infty$.

de Broglie Wavelength ($\lambda = \frac{h}{mv}$)

The wavelength associated with any piece of matter of mass $m$ moving at velocity $v$, demonstrating that particles exhibit wave-like properties at the atomic scale.

Heisenberg Uncertainty Principle

A fundamental quantum principle stating that it is impossible to simultaneously determine both the exact position and the exact momentum (or energy) of an atomic-scale particle such as an electron.

Wave Function ($\Psi$)

A mathematical description of the state of an electron in quantum mechanics derived from the Schrödinger equation, where $\Psi^2$ represents the probability density of finding the electron in a given region of space.

Atomic Orbital

A three-dimensional region of space around the nucleus where an electron of a specific quantized energy has a high probability of being located.

s Orbital

A spherically symmetric atomic orbital corresponding to angular momentum quantum number $l = 0$, consisting of one orbital per principal energy shell.

p Orbital

A dumbbell-shaped atomic orbital with two lobes and a planar node through the nucleus, corresponding to $l = 1$ and existing in a degenerate set of three ($p_x$, $p_y$, $p_z$).

d Orbital

An atomic orbital corresponding to angular momentum quantum number $l = 2$, existing in a degenerate set of five orbitals per shell ($d_{z^2}$, $d_{x^2-y^2}$, $d_{xy}$, $d_{yz}$, $d_{zx}$).

Aufbau Principle

The rule governing electron configurations stating that electrons occupy the lowest-energy available orbitals first before populating higher-energy levels.

Hund's Rule

The rule stating that when filling degenerate orbitals, electrons enter each orbital singly with parallel spins before any orbital is doubly occupied.

Core Electrons

Electrons in completely filled inner shells (corresponding to the preceding noble gas plus any completed d subshells) that shield valence electrons and do not participate in chemical reactions.

Valence Electrons

The electrons located in an atom's outermost principal shell and incomplete subshells that dictate chemical reactivity and bond formation.

Coulomb's Law

The electrostatic law given by $F = k \frac{q_1 q_2}{r^2}$, describing the attractive force between opposite charges and repulsive force between like charges as directly proportional to charge magnitude and inversely proportional to the square of the distance.

Effective Nuclear Charge ($Z_{\text{eff}}$)

The net positive charge experienced by a valence electron from the nucleus, calculated as $Z_{\text{eff}} = Z - S$, where $Z$ is the atomic number and $S$ is the core electron screening constant.

Atomic Radius

Half the distance between the nuclei of two identical bonded atoms; it decreases across a period due to increasing $Z_{\text{eff}}$ and increases down a group due to valence electrons occupying higher principal shells ($n$).

Cation

A positively charged ion formed when an atom loses one or more valence electrons, resulting in an ionic radius smaller than that of the parent atom due to decreased electron-electron repulsions and loss of the outermost orbital.

Anion

A negatively charged ion formed when an atom gains one or more electrons into its valence shell, resulting in an ionic radius larger than the neutral atom due to increased electron-electron repulsions.

Isoelectronic Series

A group of atoms or ions possessing the identical number of electrons and electron configurations, where ionic size decreases monotonically as nuclear charge ($Z$) increases.

Ionization Energy (IE)

The minimum quantity of energy required to remove an electron from an isolated gaseous atom or ion in its ground state ($M(g) \rightarrow M^+(g) + e^-$).