Chapter 3: Electronic Structure of the Atom

Flashcards for reviewing key vocabulary related to the electronic structure of the atom based on lecture notes.

Rutherford's gold foil experiment

Led to the nuclear model of an atom.

Classical mechanics

Explained physical properties of matter on the macroscopic level but could not explain properties at the atomic level.

Quantum mechanics

Attempts to describe matter at the atomic level.

Electromagnetic radiation

A form of energy that can be thought of as an oscillating wave moving through space.

Amplitude

Vertical distance from peak or trough of a wave to its midline.

Intensity

The square of a wave's amplitude, related to its brightness.

Wavelength (λ)

The peak-to-peak or trough-to-trough distance of a wave.

Frequency (ν)

How often a complete wave cycle passes a fixed point in space.

Hertz (Hz)

The unit of frequency, equal to 1 cycle per second.

Speed of light (c)

The constant speed at which electromagnetic waves travel through a vacuum, equal to 2.998 × 10^8 m/s.

Inversely proportional (wavelength and frequency)

Describes the relationship where one quantity increases as the other decreases, as seen in c = λν.

Diffraction

The spreading out of waves as they pass through closely spaced slits.

Diffraction pattern

Bands of alternating intensity created when wave fronts collide or interfere with each other.

Constructive interference

Occurs when waves combine to produce a larger amplitude, resulting in bright bands.

Destructive interference

Occurs when waves combine to produce zero amplitude, resulting in dark bands.

Electromagnetic spectrum

The range of all types of electromagnetic radiation.

Planck's constant (h)

A fundamental physical constant relating energy and frequency, equal to 6.626 × 10^-34 J·s.

Atomic spectroscopy

The study of the specific colors of light emitted by different elements.

Photoemission spectrum

The separated component wavelengths or frequencies of light emitted by an excited gas.

Continuous spectrum

A spectrum containing all wavelengths of light without gaps.

Line spectrum

A spectrum containing only specific discrete wavelengths of light, characteristic of excited elements.

Rydberg constant (R)

A constant used in the Rydberg Equation for calculating wavelengths of radiation emitted by hydrogen.

Blackbody

An idealized object capable of absorbing and emitting all frequencies of radiation.

Blackbody radiation

The energy emitted by an idealized blackbody.

Ultraviolet catastrophe

The failure of classical mechanics to accurately predict blackbody radiation at high frequencies.

Quanta

Discrete, quantized packets of energy absorbed and emitted by a blackbody, as proposed by Max Planck.

Photoelectric effect

When a metal absorbs sufficient energy, its electrons can be ejected with measurable velocity.

Threshold frequency

The minimum frequency of light required to eject electrons from a metal in the photoelectric effect.

Photon

A discrete particle or packet of light, each having an energy E = hν, as characterized by Albert Einstein.

Quantized electronic energy

The concept that the electronic energy in an atom can only exist at specific discrete values.

Bohr model of the atom

Proposes that electrons travel in circular paths around the nucleus in stationary orbits.

Excited electron

An electron that has absorbed incident energy and been promoted to a higher energy state.

Ground state

The lowest possible energy state for the hydrogen atom, where n = 1.

Excited states

Any energy states for an electron where n > 1.

Wave-particle duality of matter

The concept that all matter, especially subatomic particles, can be treated as a moving wave.

de Broglie wavelength

The wavelength associated with a moving particle (λ = h/mv).

Heisenberg Uncertainty Principle

It is impossible to describe with absolute certainty both the exact position and momentum of an electron simultaneously.

Wavefunction (Ψ)

A mathematical description of a quantum particle, where |Ψ|^2 is the probability density of the electron.

Quantum numbers

Numbers that quantize certain properties of the electron, derived from Schrödinger’s equation.

Principal quantum number (n)

Quantizes the energy of an electron and determines its most likely distance from the nucleus, defining the electron shell.

Orbital angular momentum quantum number (ℓ)

Quantizes the angular momentum of an electron and determines the shape of its orbital, defining the electron subshell.

spdf notation

A system to denote subshells based on their ℓ values: s (ℓ=0), p (ℓ=1), d (ℓ=2), f (ℓ=3).

Magnetic quantum number (mℓ)

Gives orbitals their specific orientations in space.

Spin quantum number (ms)

Describes the intrinsic angular momentum of an electron, having values of +1/2 ('spin up') or -1/2 ('spin down').

Electron configurations

A convenient way to communicate the quantum numbers of all electrons in an atom.

Orbital diagrams

Visual representations where horizontal lines represent orbitals and half arrows represent electrons.

Aufbau Principle

States that the lowest energy orbitals are fully filled before filling orbitals of higher energy.

Pauli exclusion principle

States that every electron in an atom must possess a completely unique set of quantum numbers, meaning an orbital can hold a maximum of two electrons with opposite spins.

Hund's rule

States that the ground-state electron configuration will consist of the placement of electrons into orbitals such that the total electron spin is maximized.

Effective nuclear charge (Zeff)

The net positive charge experienced by an electron in a multi-electron atom, accounting for the shielding effect of other electrons.

Metallic character

The tendency of an element to readily lose an electron.