Introductory Chemistry – Light, Quantum Theory, and Electron Configuration

Vocabulary flashcards covering light, electromagnetic waves, quantum theory, atomic models, and electron configuration presented in the lecture notes.

Electromagnetic Radiation

Energy that travels through space as oscillating electric and magnetic fields (light is one form).

Electromagnetic Spectrum

The full range of electromagnetic radiation, from radio waves to gamma rays, arranged by wavelength or frequency.

Wavelength (λ)

The distance between identical points on consecutive waves; usually measured in meters, nanometers (nm), or Ångströms (Å).

Frequency (ν)

The number of wave cycles that pass a point each second; measured in hertz (Hz).

Hertz (Hz)

The SI unit of frequency equal to one cycle (wave) per second.

Speed of Light (c)

A constant value of 3.00 × 10⁸ m s⁻¹ describing how fast light travels in a vacuum.

Inverse Relationship (λ–ν)

Wavelength and frequency are inversely proportional: c = λν.

Photon

A quantum (packet) of electromagnetic energy that carries energy proportional to its frequency.

Planck’s Constant (h)

A fundamental constant, 6.626 × 10⁻³⁴ J s, used to relate photon energy to frequency.

Photon Energy Equation

E = hν (or E = hc/λ) describes the energy of a single photon.

Visible Light

Portion of the spectrum humans see (≈ 400–700 nm), colors ROYGBIV.

Infrared (IR) Radiation

Electromagnetic waves with wavelengths longer than visible light (≈ 700 nm–1 mm); felt as heat.

Ultraviolet (UV) Radiation

Radiation with wavelengths shorter than visible light (≈ 10–400 nm); causes sunburn.

Microwave Radiation

Long-wavelength, low-frequency radiation used in cooking and communication (≈ 1 mm–1 m).

Radio Waves

Longest-wavelength electromagnetic waves used for broadcasting (≈ >1 m).

X-Rays

High-energy radiation (≈ 0.01–10 nm) used in medical imaging.

Gamma Rays

Highest-energy, shortest-wavelength radiation (< 0.01 nm) emitted by nuclear processes.

Line Spectrum

A series of discrete colored lines emitted or absorbed by an element, indicating specific energy transitions.

Flame Test

A qualitative test where metal ions emit characteristic colors when heated in a flame.

Photoelectric Effect

Ejection of electrons from a metal surface when light of sufficient frequency shines on it.

Bohr Model

1913 atomic model where electrons orbit the nucleus in fixed energy levels and emit/absorb light when jumping levels.

Ground State

The lowest possible energy arrangement of electrons in an atom.

Excited State

Any electron arrangement with higher energy than the ground state; electrons occupy higher orbitals.

Quantum Model

Modern atomic theory treating electrons as wave functions occupying orbitals, not fixed paths.

Heisenberg’s Uncertainty Principle

It is impossible to know simultaneously the exact position and momentum of a particle such as an electron.

Wave-Particle Duality

Concept that particles like electrons exhibit both wave and particle characteristics.

Principal Quantum Number (n)

Integer (1,2,3…) identifying an electron’s main energy level and relative distance from the nucleus.

Energy Level Capacity

Maximum electrons per level: 2, 8, 18, 32 for n =1–4 respectively.

Sublevel (s, p, d, f)

Division within an energy level characterized by shape and energy: s(1), p(3), d(5), f(7) orbitals.

Orbital

A region in an atom where there is a high probability of finding up to two electrons.

Electron Spin

Intrinsic magnetic property; paired electrons in an orbital have opposite spins.

Hund’s Rule

Electrons fill degenerate orbitals singly with parallel spins before pairing.

Electron Configuration

Notation showing distribution of electrons among orbitals (e.g., 1s²2s²2p⁶).

Valence Electrons

Electrons in the highest occupied energy level; participate in bonding.

Valence Level

The atom’s highest occupied principal energy level.

Octet Rule

Atoms gain, lose, or share electrons to achieve eight valence electrons (filled s and p sublevels).

Noble Gas Configuration

Stable electron arrangement resembling a noble gas with a filled valence level.

Noble-Gas Shorthand

Electron configuration starting with the previous noble gas in brackets followed by remaining orbitals.

Ion Electron Configuration

Electron arrangement of an ion after electrons are gained or lost to achieve stability (e.g., Na⁺ = [Ne]).

Inner (Core) Electrons

Electrons not in the valence level; they shield the nucleus from valence electrons.

Outer Electrons

Electrons beyond the noble-gas core; include valence electrons.

s-Block

Periodic-table columns where the outermost electrons occupy an s orbital (Groups 1 and 2).

p-Block

Columns where outermost electrons occupy p orbitals (Groups 13–18).

d-Block (Transition Metals)

Middle section where filling of d orbitals occurs; variable oxidation states common.

f-Block

Lanthanides and actinides where f orbitals are being filled.

Electron Configuration Order

Orbitals fill in the order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

Filled Sublevel Capacities

s:2 e⁻, p:6 e⁻, d:10 e⁻, f:14 e⁻.

Energy Transition (Hydrogen)

Specific electron jumps (e.g., 3→2) produce spectral lines: red, blue, indigo, violet.

Photon Absorption

Electron gains energy and moves to a higher level; light is absorbed.

Photon Emission

Electron loses energy and falls to a lower level; light is emitted.