Quantum Chemistry & Atomic Structure

Flashcards covering key vocabulary from the lecture on the electromagnetic spectrum, atomic structure, quantum numbers, orbital shapes, electron configurations, and periodic trends.

Electromagnetic Radiation (EMR)

Light, which behaves like a wave.

Wavelength (λ)

The distance between two consecutive peaks of a wave, measured in meters (m) or nanometers (nm).

Frequency (ν)

The number of waves that pass a point per second, measured in Hertz (Hz).

Amplitude

The height of a wave, which relates to its brightness.

Speed of Light (c)

A constant value of 3.000 x 10^8 m/s, related to wavelength and frequency by the equation c = λν.

Electromagnetic Spectrum

The range of all types of electromagnetic radiation, ordered by increasing energy from Radio waves to Gamma rays.

Continuous Spectrum

A spectrum that contains all wavelengths, such as that produced when white light passes through a prism.

Line Spectrum

A spectrum that shows only specific discrete wavelengths, providing evidence that electrons occupy specific quantized energy levels.

Quanta

Discrete packets of energy, according to Planck's concept that energy is quantized.

Energy of a Photon (E)

The energy carried by a single photon, calculated by the equation E = hν (where h is Planck's constant).

Planck's Constant (h)

A fundamental constant equal to 6.626 x 10^-34 J·s.

Photoelectric Effect

The phenomenon where a photon gives energy to an electron, and if that energy exceeds a threshold, the electron is ejected. Frequency (energy per photon) is crucial, while intensity affects the number of photons.

Bohr Model

A model proposing that electrons in hydrogen atoms occupy specific quantized orbits, with energy levels given by En = -2.178 x 10^-18 Z^2/n^2 J. It explains the hydrogen spectrum but fails for multi-electron atoms.

de Broglie Wavelength (λ)

The wavelength associated with a particle, such as an electron, calculated by the equation λ = h/mv. This concept demonstrates that particles have wave-like properties.

Principle Quantum Number (n)

Describes the main energy level of an electron; higher 'n' corresponds to a larger, higher-energy orbital.

Azimuthal (Angular Momentum) Quantum Number (l)

Describes the shape of an electron's orbital (s, p, d, f) and can take integer values from 0 to (n-1).

s orbital

A spherical orbital corresponding to l=0.

p orbital

A dumbbell-shaped orbital corresponding to l=1, with three possible orientations (px, py, pz).

d orbital

An orbital with typically four clover-shaped lobes, corresponding to l=2, with five possible orientations.

Aufbau Principle

States that electrons fill the lowest-energy orbitals first.

Pauli Exclusion Principle

States that no two electrons in an atom can have the same set of four quantum numbers (n, l, ml, ms).

Hund's Rule

States that for degenerate orbitals (same energy), electrons will first occupy each orbital singly with parallel spins before pairing up.

Atomic Radius

The size of an atom. It decreases across a period due to increased nuclear charge and increases down a group due to higher principal quantum numbers.

Ionic Size (Cations)

Cations are smaller than their neutral atoms because they have fewer electrons, leading to less electron-electron repulsion and a stronger pull from the nucleus.

Ionic Size (Anions)

Anions are larger than their neutral atoms because they have gained extra electrons, increasing electron-electron repulsion and expanding the electron cloud.

Ionization Energy (IE)

The energy required to remove an electron from a gaseous atom. It increases across a period and decreases down a group.

Photoelectron Spectroscopy (PES)

A technique that uses known photon energy to eject electrons from an atom and measures their kinetic energy, allowing determination of electron binding energies.

Electron Affinity (EA)

The energy change that occurs when a gaseous atom gains an electron. A negative EA indicates that energy is released, meaning the atom readily accepts an electron.