Atomic Theory, Quantum Mechanics, and Chemical Quantities

Vocabulary practice flashcards covering fundamental atomic theory, quantum mechanical models, quantum numbers, electron configuration notations, isotopes, and mole concepts.

Law of conservation of mass

The principle that mass is neither created nor destroyed during a chemical reaction or physical change in a closed system.

Law of definite proportions

The principle that a specific chemical compound always contains the same elements in the same fixed proportions by mass.

Law of multiple proportions

The principle that if elements form multiple compounds, the masses of $1$ element combining with a fixed mass of the other are in ratios of small whole numbers.

Electron

A subatomic particle with a negative charge.

Atomic nucleus

The dense central region of an atom composed of protons and neutrons.

Proton

A stable subatomic particle carrying a positive charge.

Neutron

A subatomic particle possessing no electric charge.

Electromagnetic spectrum

The complete range of all types of electromagnetic radiation.

Wavelength

The distance over which the shape of a wave repeats.

Frequency

The rate at which a repeating event occurs over an amount of time.

Photoelectric effect

The emission of electrons from a material when light shines upon it.

Quantum

The minimum amount of energy that can be gained or lost by an atom.

Photon

The smallest possible packet of electromagnetic energy, forming the basic unit of all light.

Ground state

The stationary state of lowest possible energy in a physical system.

Excited state

An energy level in a physical system that has a higher energy than its lowest baseline configuration.

Line-emission spectrum

A pattern of brightly colored lines on a dark background produced by gas atoms releasing light of specific wavelengths.

Bohr model

An atomic model that depicts electrons orbiting a central nucleus along defined circular paths.

Electron shell

An orbit or energy level that electrons follow around the nucleus of an atom.

Valence electrons

Electrons located in the outermost shell that participate in chemical bonding and govern chemical reactivity.

Octet rule

The principle that main-group atoms tend to lose, gain, or share electrons to attain a full outer valence shell of $8$ electrons.

Lewis Structure

A diagram illustrating how valence electrons are distributed around atoms within a molecule.

Quantum model

An atomic model describing electrons as waves and determining the probability of finding them in designated regions called orbitals.

Heisenberg uncertainty principle

The principle establishing that one cannot simultaneously know both the exact position and exact momentum of a quantum particle.

Schrodinger wave equation

A fundamental equation in quantum mechanics predicting how a physical system evolves over time and describing the wave-like properties of particles such as electrons.

Orbital

A mathematical region of space around an atomic nucleus where an electron is most likely to be found.

Quantum numbers

A set of $4$ numerical values that designate a unique energy state and spatial address for an electron in an atom.

Principal quantum number ($n$)

One of the $4$ quantum numbers used in physics and chemistry to indicate the main energy level and state of an electron in an atom.

Angular momentum quantum number ($l$)

The quantum number determining the shape of an atomic orbital and the magnitude of its orbital angular momentum.

Magnetic quantum number ($m$)

The quantum number describing the specific three-dimensional spatial orientation of an atomic orbital.

Spin quantum number ($m_s$)

An intrinsic property describing the angular momentum and direction of spin of a subatomic particle such as an electron.

Electron configuration

The arrangement and distribution of an atom's or molecule's electrons among atomic or molecular orbitals.

Aufbau principle

The rule dictating that electrons populate an atom's lowest available energy orbitals before occupying higher-energy orbitals.

Pauli exclusion principle

The principle stating that no $2$ identical particles with half-integer spin can occupy the identical quantum state simultaneously.

Hund's rule

The rule stating that every orbital in a subshell is singly occupied with $1$ electron before any orbital is doubly occupied, and all single electrons share parallel spins.

Atomic number

The total number of protons found in the nucleus of an atom.

Mass number

The total sum of protons and neutrons located within an atomic nucleus ($\text{mass number} = \text{protons} + \text{neutrons}$).

Isotopes

Atoms belonging to the same chemical element that possess the identical number of protons but varying numbers of neutrons.

Average atomic mass

The weighted average mass of an element's naturally occurring isotopes, measured in units of $\text{amu}$.

Noble Gas notation

An abbreviated method of writing electron configurations that uses the bracketed symbol of the preceding noble gas (e.g., $[\text{He}]$) to represent core electrons.

Orbital notation

A visual representation showing the subshells and individual orbitals of an atom with upward and downward arrows indicating opposite electron spins.

Nuclear symbol

A symbolic format showing the element symbol along with its mass number as a superscript and atomic number as a subscript (e.g., $^{12}_{6}\text{C}$, $^{13}_{6}\text{C}$, $^{14}_{6}\text{C}$).

Hyphen notation

A notation style specifying an isotope by writing the element name followed by a hyphen and its mass number (e.g., uranium-$235$, uranium-$238$, boron-$10$).

Mole

A fundamental scientific unit of measurement representing an amount of a substance.

Molar mass

The total mass corresponding to $1\text{ mole}$ of a given chemical substance.

Avogadro's number

The number of constituent units in $1\text{ mole}$ of a substance, given as $6.02214086 \times 10^{23}$.

John Dalton

Scientist that discovered the atom; said atoms are solid spheres

Joseph Thomson

Discovered electrons using cathode rays; Plum Pudding Model

Ernest Rutherford

Discovered the nucleus + protons using the Gold Foil Experiment

Albert Einstein and Max Planck

Discovered the dual wave-particle nature of light

Niels Bohr

Hypothesized that electrons orbit around the nucleus and exist in quantized energies

Louis de Brogile

Discovered that electrons have a dual wave-particle nature

Werner Heisenberg

Discovered that it’s impossible to know the location of electrons at a given time

Erwin Schrödinger

Made the quantum model and discovered that electrons occupy 3D orbitals