Honors Chemistry: Atomic Theory, Light, and Quantum Mechanics

Comprehensive vocabulary flashcards covering the properties of light, historical atomic theory, quantum numbers, and electron configuration principles.

Electromagnetic Radiation

A form of energy that exhibits wave-particle duality and travels through space at the speed of light ($c = 3.00 \times 10^8\,\text{m/s}$ in a vacuum).

Primary Colors of Light

Red, green, and blue light, which combine in equal amounts to produce white light and can combine in various amounts to form any color.

Secondary Colors of Light

Colors produced when two primary colors combine in equal amounts, consisting of magenta, yellow, and cyan.

Wavelength ($\lambda$)

The distance between corresponding points on adjacent waves, such as crest to crest or trough to trough.

Frequency ($\nu$)

The number of complete waves that pass a given point in one second, measured in Hertz ($\text{Hz}$) or reciprocal seconds ($\text{s}^{-1}$).

Quantum

The minimum amount of energy that can be gained or lost by an atom, introduced by Max Planck in his study of light emission by hot objects.

Planck's Constant ($h$)

A fundamental constant with a value of $6.626 \times 10^{-34}\,\text{J}\cdot\text{s}$ that relates the energy of a quantum to its frequency ($E = h\nu$).

Photon

A discrete particle of electromagnetic radiation proposed by Albert Einstein in 1905 that carries a quantum of energy ($E = h\nu$) and has zero rest mass.

Photoelectric Effect

The phenomenon in which electrons are ejected from the surface of a metal when exposed to light of a sufficiently high minimum frequency.

Ground State

The lowest potential energy state of an electron in an atom.

Excited State

A state in which an electron has absorbed energy and attained a higher potential energy than its ground state.

Line-Emission Spectrum

The distinct series of specific frequencies and wavelengths of light expelled when excited electrons fall to lower energy levels (such as the four colors produced by hydrogen).

Bohr Model

Niels Bohr's 1913 model of the hydrogen atom describing electrons circling the nucleus in set orbits, where orbital distance correlates directly with electron energy.

Atom (Democritus)

Nature's basic indivisible particle, named from the Greek word for 'indivisible' around 400 BC.

Law of Conservation of Mass

A fundamental law established by Antoine Lavoisier stating that mass is neither created nor destroyed during ordinary chemical reactions.

Law of Definite Proportions

Joseph Proust's law stating that a specific chemical compound always contains the exact same elements in identical proportions by mass regardless of sample size or source.

Law of Multiple Proportions

John Dalton's 1803 law stating that when two elements combine to form more than one compound, the mass ratios of the second element combining with a fixed mass of the first are small whole numbers.

Dalton's Atomic Theory

The 1803 theory stating that matter consists of indivisible atoms; atoms of a given element are identical in mass and properties; and atoms combine, separate, or rearrange in chemical reactions.

Electron

A negatively charged subatomic particle discovered in 1897 by J.J. Thomson through experiments observing the deflection of cathode rays.

Plum Pudding Model

J.J. Thomson's model representing an electrically neutral atom as a uniform positive sphere with electrons embedded evenly throughout.

Rutherford Gold Foil Experiment

Ernest Rutherford's 1911 experiment bombarding gold foil with alpha particles, where 1 in 8,000 deflections demonstrated the existence of a small, dense, positively charged nucleus.

Nucleus

The dense, positively charged center of an atom that contains almost all of its mass and leaves the remainder of the atom as mostly empty space.

Neutron

A neutral subatomic particle located in the atomic nucleus, discovered by James Chadwick in 1932.

De Broglie Hypothesis

Louis de Broglie's 1924 proposal that electrons exhibit wave-particle duality, acting as confined waves of specific frequencies around the nucleus.

Heisenberg Uncertainty Principle

Werner Heisenberg's 1927 principle stating that it is fundamentally impossible to simultaneously determine both the exact position and velocity of an electron.

Schrödinger Wave Equation

An equation formulated by Erwin Schrödinger in 1926 treating electrons mathematically as waves, establishing the foundation of modern quantum theory.

Orbital

A three-dimensional region around the nucleus representing the high-probability location of an electron, which can hold a maximum of two electrons.

Principal Quantum Number ($n$)

The quantum number indicating the main energy level of an electron, where the total number of orbitals in that energy level equals $n^2$.

Angular Momentum Quantum Number ($l$)

The quantum number indicating orbital shape or sublevel, taking values from $0$ to $n - 1$ ($l = 0$ is spherical $s$, $l = 1$ is dumbbell $p$, $l = 2$ is complex $d$, $l = 3$ is complex $f$).

Magnetic Quantum Number ($m$)

The quantum number designating the spatial orientation of an orbital around the nucleus, taking integer values from $-l$ to $+l$.

Spin Quantum Number

The quantum number indicating the intrinsic spin orientation of an electron, which can only equal $+\frac{1}{2}$ or $-\frac{1}{2}$.

Pauli Exclusion Principle

The quantum mechanical principle stating that no two electrons in the same atom can possess identical sets of all four quantum numbers, requiring paired electrons to have opposite spins.

Aufbau Principle

The structural rule stating that an electron occupies the lowest-energy orbital available that can receive it before higher-energy levels are filled.

Diagonal Rule

A mnemonic visual pattern of diagonal arrows used to determine the exact order of sublevel orbital filling by increasing energy according to the Aufbau principle.

Hund's Rule

The rule stating that degenerate orbitals of equal energy are each occupied by one electron with parallel spin before any orbital receives a second pairing electron.

Electron Configuration

The formal notation describing the ground-state arrangement of all electrons in an atom across primary energy levels and sublevels (e.g., Nitrogen is $1s^2 2s^2 2p^3$).