Atomic Structure, Light, and Periodic Trends

Vocabulary flashcards covering light, spectra, wave-particle duality, quantum numbers, electronic rules, shielding, effective nuclear charge, and periodic trends.

Atomic emission line

A bright line produced when an electron in an excited state drops from a higher energy level to a lower energy level, releasing the energy difference ($\Delta E$) as a single photon of light ($E = h\nu$).

Uniqueness of atomic spectra

The distinct set of quantized energy levels, $\Delta E$ values, and emitted or absorbed photon wavelengths specific to each element due to its unique number of protons and electrons creating unique electrostatic forces.

Absorption spectrum

A spectrum appearing as dark lines on a continuous rainbow background, formed when white light passes through a cold gas and atoms absorb photons whose energy matches the exact $\Delta E$ required to jump from lower to higher energy levels.

Wave nature of light (CER)

Claim: Light behaves as a wave. Evidence: Light passing through a double slit creates an interference pattern of bright and dark fringes. Reasoning: Waves interfere constructively (in-phase) and destructively (out-of-phase), a property unique to waves.

Particle nature of light (CER)

Claim: Light behaves as a particle (photon). Evidence: The photoelectric effect shows a threshold frequency below which no electrons are ejected, regardless of brightness. Reasoning: Light energy comes in discrete packets (photons); a single photon must have enough energy to eject a single electron.

Electromagnetic spectrum order

The arrangement of light from lowest energy/longest wavelength to highest energy/shortest wavelength: Radio waves $\rightarrow$ Microwaves $\rightarrow$ Infrared (IR) $\rightarrow$ Visible $\rightarrow$ Ultraviolet (UV) $\rightarrow$ X-rays $\rightarrow$ Gamma rays.

Light intensity (photoelectric effect)

An increase in light brightness above the threshold frequency that causes more photons to strike the metal per second, ejecting a greater number of electrons (increasing current) while individual electrons leave with the exact same kinetic energy.

Heisenberg Uncertainty Principle

A principle stating that it is impossible to simultaneously know both the exact position and momentum (velocity) of a small particle like an electron, thereby invalidating Bohr's model of fixed circular orbits.

Principal quantum number ($n$)

The quantum number representing the main energy level and shell size of an electron.

Angular momentum quantum number ($l$)

The quantum number representing subshell shape, where $s = 0$, $p = 1$, $d = 2$, and $f = 3$.

Magnetic quantum number ($m_l$)

The quantum number indicating the spatial orientation of an orbital in three-dimensional space.

Electron spin quantum number ($m_s$)

The quantum number representing the spin direction of an electron, which can have values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

$f$ subshell capacity

A subshell that contains $7$ orbitals and can hold a maximum of $14$ electrons.

Aufbau Principle

The quantum mechanical rule stating that electrons fill lower-energy atomic orbitals before filling higher-energy ones.

Hund's Rule

The rule stating that when filling degenerate (equal energy) orbitals, electrons occupy them singly with parallel spins before pairing up.

Pauli Exclusion Principle

The principle stating that no two electrons in the same atom can have the exact same set of four quantum numbers; an orbital can hold a maximum of two electrons with opposing spins.

Shielding

The process by which core electrons block valence electrons from experiencing the full electrostatic pull of the nucleus.

Effective Nuclear Charge ($Z_{eff}$)

The net positive charge experienced by valence electrons, calculated using the formula $Z_{eff} = Z - S$, where $Z$ is the number of protons and $S$ is the number of core electrons.

Atomic radius group trend

The trend where atomic size increases moving down a group because additional principal energy levels ($n$) are added, placing valence electrons further from the nucleus (larger $r$).

Atomic radius periodic trend

The trend where atomic size decreases across a row (left to right) because protons are added while core shielding remains constant, causing $Z_{eff}$ to increase and pull the valence electron cloud closer.

Cation and anion radius trends

Cations (+) are smaller than their neutral parent atoms because losing electrons decreases electron-electron repulsion and can empty an outer shell; anions (-) are larger because adding electrons increases electron-electron repulsion, expanding the electron cloud.

Isoelectronic series radius rule

In a series of chemical species with the exact same number of electrons, the species with the most protons has the smallest radius because the greater nuclear charge pulls the electron cloud inward more tightly.

Ionization Energy (IE)

The energy required to remove an electron from an atom in the gas phase.