Chapter 3: Atomic Structure - Explaining the Properties of Elements

Comprehensive vocabulary flashcards covering atomic structure, electromagnetic radiation, quantum mechanics, quantum numbers, orbital filling rules, electron configurations, and periodic trends from Chapter 3 of Chemistry: An Atoms-Focused Approach.

Electromagnetic Radiation

Any form of radiant energy in the electromagnetic spectrum that propagates through space.

Electromagnetic Spectrum

A continuous range of radiant energy that includes gamma rays, X-rays, ultraviolet radiation, visible light, infrared radiation, microwaves, and radio waves.

Speed of Light ($c$)

The constant speed of light in a vacuum ($2.998 \times 10^8\,\text{m/s}$), which relates wavelength ($\lambda$) and frequency ($\nu$) through the equation $\lambda \nu = c$.

Fraunhofer Lines

A series of narrow dark lines observed within the spectrum of sunlight, which match the characteristic atomic emission spectra of specific elements.

Atomic Emission Spectra

Characteristic patterns of bright lines produced when atoms are vaporized and excited in high-temperature flames or electrical discharges.

Atomic Absorption Spectra

Characteristic patterns of dark lines produced when an external source of radiation passes through a sample of free gaseous atoms.

Blackbody Radiator

An ideal source of radiant energy that is a perfect absorber of light when cold and a perfect radiator of light when hot.

Quantum

The smallest discrete quantity of a particular form of energy.

Planck Constant ($h$)

A fundamental physical constant with a value of $6.626 \times 10^{-34}\,\text{J}\cdot\text{s}$, relating the frequency of radiation to its energy via $E = h\nu$.

Quantized

Having values restricted to whole-number multiples of a specific base value, analogous to discrete steps on a staircase rather than a continuous ramp.

Photon

A tiny packet of radiant energy; a single quantum of electromagnetic radiation with energy calculated as $E = h\nu = \frac{hc}{\lambda}$.

Photoelectric Effect

The release or ejection of electrons from a material or metal surface when it absorbs incident electromagnetic radiation.

Threshold Frequency ($\nu_0$)

The minimum frequency of incident light required to dislodge and eject an electron from the surface of a metal.

Work Function ($\Phi$)

The minimum amount of energy required to dislodge an electron from the surface of a metal, given by $\Phi = h\nu_0$.

Ground State

The most stable, lowest-energy state of an atom or particle, corresponding to $n = 1$ in the hydrogen atom.

Excited State

Any allowed energy state of an atom or particle that is higher in energy than its ground state.

Lyman Series

A series of spectral emission transitions in the hydrogen atom occurring in the ultraviolet region, where excited electrons drop down to the $n = 1$ ground state.

Balmer Series

A series of spectral emission transitions in the hydrogen atom occurring in the visible region, where excited electrons relax down to the $n = 2$ level.

Paschen Series

A series of spectral emission transitions in the hydrogen atom occurring in the infrared region, where excited electrons relax down to the $n = 3$ level.

Matter Waves

The wave-like behavior exhibited by matter particles such as electrons, proposed by Louis de Broglie with wavelength given by $\lambda = \frac{h}{mu}$.

Heisenberg Uncertainty Principle

The fundamental quantum principle stating that one cannot simultaneously determine the exact position and exact momentum/velocity of an electron: $\Delta x \cdot m\Delta u \ge \frac{h}{4\pi}$.

Wave Function ($\psi$)

A mathematical solution to the Schrödinger wave equation that describes how an electron's matter wave varies with location inside an atom.

Atomic Orbital

A three-dimensional region in space around an atom's nucleus where there is a high probability of finding an electron, defined mathematically by $\psi^2$.

Principal Quantum Number ($n$)

A positive integer ($n = 1, 2, 3, \dots$) that indicates the relative size and main energy level of an orbital (principal shell).

Angular Momentum Quantum Number ($\ell$)

An integer ranging from $0$ to $(n - 1)$ that defines the three-dimensional geometric shape of an orbital (subshell letters $s, p, d, f$).

Magnetic Quantum Number ($m_\ell$)

An integer with values ranging from $-\ell$ to $+\ell$ that defines the spatial orientation of an orbital in three dimensions.

Spin Quantum Number ($m_s$)

A quantum number taking values of $+\frac{1}{2}$ or $-\frac{1}{2}$ that describes the intrinsic spin orientation of an electron.

Aufbau Principle

The rule for determining ground-state electron configurations by sequentially placing electrons into the lowest-energy available orbitals first.

Pauli Exclusion Principle

The principle stating that no two electrons in an atom can possess the exact same set of four quantum numbers ($n, \ell, m_\ell, m_s$); thus, each orbital can hold at most two electrons with opposite spins.

Hund's Rule

The principle that the lowest-energy electron configuration in degenerate orbitals is the one containing the maximum number of unpaired electrons with parallel (identical) spins.

Degenerate Orbitals

Orbitals within the same subshell that have identical energy levels (such as the three separate $2p$ orbitals in a multielectron atom).

Core Electrons

Electrons located in the completely filled, inner shells of an atom or ion that do not participate in chemical bonding or reactions.

Valence Electrons

Electrons residing in the outermost occupied shell of an atom that exert the greatest influence on its chemical reactivity and bonding.

Valence Shell

The outermost occupied principal energy shell ($n$) of an atom.

Orbital Penetration

The ability of an electron in an outer orbital to be found close to the nucleus, causing the relative subshell energies within a shell to follow the order $s < p < d < f$.

Isoelectronic

A term describing atoms or ions that possess the identical ground-state electron configuration (such as $\text{Na}^+$ and $\text{F}^-$, both having the configuration of $\text{Ne}$).

Effective Nuclear Charge ($Z_{\text{eff}}$)

The net attractive positive charge experienced by valence electrons from nuclear protons, reduced from the actual nuclear charge ($Z$) by shielding from inner electrons.

Shielding Effect

The reduction in electrostatic attraction between nuclear protons and valence electrons caused by the repulsive forces of partially or fully filled inner electron shells.

Atomic Radius

A measure of atomic size that increases down a column due to higher principal quantum numbers ($n$) and decreases across a period due to increasing effective nuclear charge ($Z_{\text{eff}}$).

Electronegativity

The relative capacity of an atom in a chemical bond to attract shared bonding electrons toward itself, commonly measured using the Pauling scale.

First Ionization Energy ($\text{IE}_1$)

The quantity of energy required to remove one mole of electrons from one mole of isolated gaseous atoms in their ground state.

Electron Affinity

The energy change that occurs when one mole of electrons combines with one mole of gas-phase atoms or ions (e.g., $\text{Cl}(g) + e^- \rightarrow \text{Cl}^-(g)$).