Electronic Structure of the Atom and Atomic Properties

Vocabulary-style flashcards covering the nature of light, quantum mechanics, electronic configurations, and periodic trends based on the lecture transcript.

Electromagnetic Radiation (EMR)

Energy propagated through space in the form of periodic oscillations of electric and magnetic fields, exhibiting a dual nature of waves and particles (photons).

Wavelength ($\lambda$)

The distance between two successive peaks or troughs in a wave, typically measured in units like $\text{m}$, $\text{nm}$ ($10^{-9}\text{m}$), or $\mu\text{m}$ ($10^{-6}\text{m}$).

Amplitude ($A$)

The maximum height of the wave above the center line or the maximum depth below it.

Frequency ($\nu$)

The number of oscillations per second, measured in units of $\text{cycle/sec}$, $\text{s}^{-1}$, or $\text{Hertz (Hz)}$.

Speed of Light ($c$)

The distance travelled by a wave per unit time; in a vacuum, EMR has a constant velocity of $3 \times 10^{8}\text{m/s}$, defined by the formula $c = \nu \times \text{λ}$.

Photons

Particles of light with zero rest mass and no electric charge, where energy is defined by $E = h\nu$.

Blackbody Radiation

Radiation covering a wide range of wavelengths emitted by heated objects, where the emitted color depends on the temperature ($T$).

Ultraviolet Catastrophe

A failure of classical physics where intensity was predicted to increase without limit at shorter wavelengths (UV region) rather than showing a maximum intensity.

Planck’s Quantum Theory

The theory stating that energy of electromagnetic waves is quantized rather than continuous, meaning energy is absorbed or emitted in discrete packets called quanta.

Planck's Constant ($h$)

A fundamental constant used to calculate the energy of a quantum, equal to $6.626 \times 10^{-34}\text{J}•\text{s}$.

Photoelectric Effect

A phenomenon where electrons are ejected from the surface of a metal when exposed to light of a frequency higher than a specific threshold ($\nu > \nu_{o}$).

Threshold Frequency ($\nu_{o}$)

The minimum frequency of light required to eject an electron from a metal surface.

Work Function ($E_{o}$)

The particular electrostatic attraction or binding energy a metal has for its electrons ($E_{o} = h\nu_{o}$) that must be overcome for emission.

Continuous Spectrum

A band of colors observed when sunlight is passed through a prism.

Atomic/Line Spectrum

A series of narrow, discrete lines observed when light emitted by atoms (like Hydrogen) passes through a prism, acting as a spectral fingerprint for an element.

Bohr Model of the H Atom

A model where electrons move in restricted orbits of fixed energy ($E_{n}$) and radius around the nucleus, defined by the principal quantum number $n$.

Rydberg’s Constant ($R_{H}$)

The constant used to calculate the energy of an electron in a Hydrogen orbit, equal to $2.179 \times 10^{-18}\text{J}$.

Ground State

The lowest possible energy state for an electron in an atom (for Hydrogen, $n=1$).

Excited State

A state where an atom's electron has moved to a higher energy level (larger $n$) by absorbing energy.

de Broglie’s Wave-Particle Duality

The concept that particles like electrons possess wave properties, related by the equation $\text{λ} = \frac{h}{mv}$.

Heisenberg’s Uncertainty Principle

The principle stating it is impossible to simultaneously know both the momentum ($p$) and the position ($x$) of a moving particle with great certainty.

Wave Function ($ψ$)

A mathematical equation from the Quantum Mechanical Model that describes the energy and probability of finding an electron in an atom.

Electron Density ($ψ^{2}$)

The square of the wave function, which gives the probability of where an electron is likely to be at any given time.

Atomic Orbital

A 3D probability density region, often enclosing 90% of electron density, where an electron is likely to be found.

Node

A zone within an atomic orbital where there is a 0% probability of finding an electron.

Principal Quantum Number ($n$)

Describes the main energy level or shell and is related to the size and average distance of an electron from the nucleus ($n = 1, 2, 3\text{...}$).

Angular/Azimuthal Quantum Number ($ℓ$)

Relates to the shape of the orbital and divides shells into subshells ($s, p, d, f, g$), with values from $0$ to $n-1$.

Magnetic Quantum Number ($m_{ℓ}$)

Describes the spatial orientation of an orbital relative to others, with values ranging from $-ℓ$ to $+ℓ$.

Spin Quantum Number ($m_{s}$)

Describes the clockwise ($+1/2$) or counterclockwise ($-1/2$) spinning motion of an electron in an orbital.

Pauli’s Exclusion Principle

States that each orbital can contain a maximum of 2 electrons and no two electrons in the same atom can have the exact same set of four quantum numbers.

Aufbau Principle

The 'building-up' principle stating that orbitals are filled with electrons in order of increasing energy, often determined by the $(n + ℓ)$ rule.

Hund’s Rule of Multiplicity

States that all degenerate orbitals must be singly occupied by electrons with parallel spins before pairing begins.

Valence Electrons

Electrons located in the outermost shells of an atom.

Diamagnetic

Describes an atom where all electrons are paired; it is weakly repelled by a magnet.

Paramagnetic

Describes an atom that has unpaired electrons; it is weakly attracted to a magnet.

Isoelectronic

Refers to different atoms or ions that have the same electronic configuration.

Effective Nuclear Charge ($Z_{eff}$)

The actual positive charge felt by an electron, calculated as $Z_{eff} = Z - σ$, where $σ$ is the shielding constant.

Shielding Effect

The process by which core electrons reduce the electrostatic attraction between the nucleus and valence electrons.

Ionization Energy ($IE$)

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

Electron Affinity ($EA$)

A thermodynamic quantity measuring the energy released when a gaseous atom gains an electron to form an anion.