Unit 2: Atomic Structure, Quantum Mechanics, and Chemical Bonding

Vocabulary practice flashcards covering atomic history, wave-particle duality, quantum numbers, electron configurations, magnetism, orbital mechanics, bonding types, and crystal structures.

Millikan's Electron Mass Measurement

The experimental measurement finding the mass of the electron to be $9.11 \times 10^{-31}\,\text{kg}$.

Radioactivity

The spontaneous decay of the nucleus of an atom.

Gold foil experiment

Rutherford's experiment showing alpha particles, beta particles, and gamma rays, which concluded that atoms are mostly empty space and contain a central positive nucleus (proton).

Photoelectric Effect

The phenomenon where electrons are emitted by matter that absorbs energy from shortwave electromagnetic radiation (such as visible light or UV rays).

Planck's Quantum Hypothesis

The theory that because atoms oscillate, energy release occurs in packets or quanta, where hot solid bodies emit energy in bursts rather than continuously.

Photon

A unit of light energy theorized by Einstein as streams of particles that deliver enough energy to eject electrons.

Threshold frequency

The minimum frequency of radiation needed to eject an electron from a material.

Principal Quantum Number ($n$)

The quantum number (also referred to as the shell or orbital number) designating the main energy level of an electron.

Secondary Quantum Number ($l$)

The quantum number ($l = n - 1$) that describes the shape of an orbital: $s$ (sharp, spherical), $p$ (principal, dumbbell), $d$ (diffuse, donut), and $f$ (fundamental, flower).

Magnetic Quantum Number ($m_l$)

The quantum number specifying the spatial orientation of an orbital in three dimensions.

Spin Quantum Number ($m_s$)

The quantum number representing the intrinsic spin orientation of an electron, taking values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

Pauli exclusion principle

The principle stating that no two electrons in the same atom can have the exact same set of four quantum numbers ($n, l, m_l, m_s$); an orbital can hold only two electrons, which must have opposite spins.

Penetration effect

The effect in multi-electron atoms where an $s$ orbital electron is closer to the nucleus for a brief moment, giving it a lower energy level such that $E_{ns} < E_{np} < E_{nd} < E_{nf}$.

Aufbau principle

The principle that electrons fill atomic orbitals starting from the lowest available energy orbital before occupying higher energy orbitals.

Hund's rule

The rule stating that within a set of orbitals of equal energy, the lowest energy electron configuration is achieved by maximizing the number of unpaired electrons with parallel spins.

Ferromagnetism

The form of magnetism found in iron (Fe), nickel (Ni), and cobalt (Co), in which small, dense domains of atoms align their magnetic moments in the presence of an external magnetic field.

Paramagnetism

The weak attraction of a material to a magnetic field, caused by the presence of unpaired electrons (seen in metals like aluminum and platinum).

Line spectrum

A spectrum exhibiting discrete, characteristic wavelengths emitted when excited electrons in an element drop to lower energy states, emitting individual photons.

Transition

The movement of an electron from one discrete energy level to another within an atom.

Ground state

The lowest possible energy state for an atom or electron.

Quantum (wave) mechanics

The application of quantum theory to describe and explain the wave-particle properties of matter, particularly electrons in atoms.

Orbital

A three-dimensional region around the nucleus where an electron has a high probability of being found.

Heisenberg uncertainty principle

The principle stating that it is fundamentally impossible to simultaneously know both the exact position and the exact speed (momentum) of an electron.

Wave function

A mathematical function that describes the quantum state of an electron in an atom and relates to the probability of finding it at a given energy.

Electron probability density

The probability of finding an electron at a specific point in space, plotted from wave equations to determine the shapes of atomic orbitals.

Coordinate Covalent Bonding

A type of covalent bond in which both bonding electrons shared between two atoms originate entirely from only one of the participating atoms (e.g., in ammonium).

Pi bond ($\pi$ bond)

A chemical bond produced by the sideways overlap of parallel unhybridized $p$ orbitals merging above and below the axis of a sigma bond.

Metallic crystals

A solid consisting of closely packed metal cations held together by electrostatic forces of attraction to a delocalized sea of free-moving electrons.

Electron Sea Theory

The model of metallic bonding describing a solid where valence electrons move freely throughout a rigid lattice of positively charged nuclei.

Molecular crystals

Solids composed of discrete molecules held together by intermolecular forces of attraction, characterized by low melting points and poor electrical conductivity.

Covalent crystal network

A crystalline solid where atoms are interconnected in a continuous network of covalent bonds, resulting in very high melting points, insolubility, and non-conductivity.

Semiconductor

A substance with low electrical conductivity at room temperature that exhibits significantly increased conductivity at higher temperatures.