Module 1: Electronic Structure of the Atom

Vocabulary flashcards covering fundamental concepts in atomic structure, quantum mechanics, quantum numbers, electronic configuration, and periodic trends based on lecture slides.

Wave

A vibrating disturbance by which energy is transmitted.

Electromagnetic Radiation

An emission or transmission of energy in the form of electromagnetic waves.

Interference

The net effect of the combination of two or more waves moving on intersecting or coincident paths.

Diffraction

A phenomenon resulting from interference where waves spread around an obstacle or through a gap.

Photoelectric Effect

A phenomenon in which electrons are ejected from the surface of certain metals exposed to light of at least a minimum threshold frequency ($\nu_0$).

Threshold Frequency ($\nu_0$)

The minimum frequency of light required to remove an electron from a metal surface in the photoelectric effect.

Dual Nature of Light

The principle that electromagnetic radiation exhibits both wave properties and particulate properties (photons).

de Broglie Equation

The formula $\lambda = \frac{h}{mv}$ which describes the wave properties of matter by relating wavelength to mass and velocity.

Heisenberg’s Uncertainty Principle

The principle stating that it is impossible to know simultaneously both the momentum ($p$) and the position ($x$) of a particle with certainty, expressed as $\Delta x \Delta p \ge \frac{h}{4\pi}$.

Wave Function ($\Psi$)

A mathematical function of the coordinates of an electron's position in three-dimensional space that satisfies the Schrödinger equation ($\hat{H}\Psi = E\Psi$).

Probability Density ($\Psi^2$)

The product of the square of the wave function, representing the total probability of finding a particle in a small volume of space.

Orbital

A wave function that is a solution to the Schrödinger equation, representing the probability distribution of finding an electron near a particular point in space.

Nodes

Regions in an atom or orbital where there is zero electron density or zero probability of finding an electron.

Principal Quantum Number ($n$)

A quantum number with positive non-zero integral values ($1, 2, 3\dots$) that relates to the size, energy, and average distance of an electron from the nucleus.

Angular Momentum Quantum Number ($\ell$)

A quantum number with integral values from $0$ to $n - 1$ that determines the shape of an orbital and defines its subshell ($s, p, d, f$).

Magnetic Quantum Number ($m_\ell$)

A quantum number with integral values from $-\ell$ to $+\ell$ (including zero) that determines the orientation of an orbital in space relative to other orbitals.

Electron Spin Quantum Number ($m_s$)

A quantum number with values of $+\frac{1}{2}$ or $-\frac{1}{2}$ that describes the intrinsic magnetic field orientation of an electron occupying an orbital.

Pauli’s Exclusion Principle

The principle stating that in any given atom, no two electrons can have the exact same set of four quantum numbers, meaning an orbital can hold a maximum of two electrons with opposite spins.

Aufbau Principle

The principle stating that as protons are added one by one to the nucleus to build up elements, electrons are added to the lowest-energy orbitals first.

Hund’s Rule of Multiplicity

The rule stating that the most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins.

Valence Electrons

The electrons located in the outermost principal quantum level ($n$) of an atom.

Paramagnetic

A property of a substance containing net unpaired electron spins, causing it to be attracted by a magnet.

Diamagnetic

A property of a substance containing no net unpaired electron spins, causing it to be slightly repelled by a magnet.

Atomic Radius

One-half the distance between the two nuclei in two adjacent metal atoms (metallic radius) or in a diatomic molecule (covalent radius).

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

The nuclear charge felt by an electron when taking into account both the actual nuclear charge ($Z$) and shielding/repulsive effects of other electrons ($\sigma$), defined as $Z_{\text{eff}} = Z - \sigma$.

Ionization Energy

The minimum energy required to remove an electron from a gaseous atom in its ground state.

Electron Affinity

The energy change associated with the addition of an electron to a gaseous atom.

Electronegativity

The tendency of an atom to pull or attract electrons toward itself in a covalent bond.

Thomson's Plum Pudding Model

An early atomic model proposed by Joseph John Thomson after discovering the electron ($e/m = -1.76 \times 10^8\,C/g$), depicting electrons embedded within a diffuse positive charge.

Rutherford's Nuclear Model

An atomic model established by Ernest Rutherford following gold foil alpha particle experiments, demonstrating that an atom's positive charge is concentrated in a tiny central nucleus while electrons orbit through empty space.