Chapter 3: Electronic Structure and Periodic Properties of Elements - Vocabulary

Vocabulary-style flashcards covering key concepts from Chapter 3: Electronic Structure and Periodic Properties of Elements.

Electromagnetic spectrum

The range of all types of electromagnetic radiation, ordered by frequency or wavelength.

Wave-particle duality

The concept that light and matter exhibit both wave-like and particle-like properties.

Wavelength (λ)

Distance between consecutive peaks or troughs of a wave.

Frequency (ν)

Number of wave cycles passing a point per unit time; units are s−1 (hertz).

Speed of light (c)

Constant speed of light in vacuum; c = 2.998 × 10^8 m/s; relates wavelength and frequency via c = λν.

Planck's constant (h)

Constant relating energy to frequency: E = hν; h = 6.626 × 10^−34 J·s.

Photon

Quantum of light with energy E = hν; exhibits particle-like behavior.

Blackbody radiation

Idealized radiation from a perfect emitter used to model spectral distribution as a function of temperature.

Ultraviolet catastrophe

Classical prediction of blackbody radiation failing at short wavelengths; resolved by quantum theory (Planck).

Bohr model

Hydrogen atom model with quantized energy levels; electrons transition between levels by absorbing/emitting photons.

Rydberg equation

Relation for hydrogen spectral lines: 1/λ = R∞(1/n1^2 − 1/n2^2).

Rydberg constant

R∞ ≈ 1.097 × 10^7 m^−1; used in hydrogen spectral predictions.

De Broglie wavelength

Matter wave concept: λ = h/p, where p is momentum; particles exhibit wave-like behavior.

Schrödinger equation

Quantum mechanical equation describing how the quantum state evolves; Ĥψ = Eψ.

Wavefunction (ψ)

Mathematical function whose magnitude squared, |ψ|^2, gives the probability density of finding a particle.

Heisenberg uncertainty principle

Δx Δp ≥ ħ/2; cannot simultaneously know exact position and momentum.

Quantum numbers

Four numbers (n, l, ml, ms) that specify the state of an electron in an atom.

Principal quantum number (n)

1, 2, 3, …; labels energy levels; energy generally increases with n.

Angular momentum quantum number (l)

Defines orbital shape; l = 0(s), 1(p), 2(d), 3(f); subshells within a shell.

Magnetic quantum number (m_l)

Orientation of the orbital in space; m_l ranges from −l to +l (2l+1 orbitals).

Spin quantum number (m_s)

Intrinsic electron spin; m_s = +1/2 or −1/2.

Pauli exclusion principle

No two electrons in an atom can have the same set of all four quantum numbers.

Hund's rule

For degenerate orbitals, the lowest-energy configuration has the maximum number of unpaired electrons.

Electron configuration

Arrangement of electrons in orbitals; constructed by the Aufbau principle and Hund's rule.

Valence electrons

Electrons in the outermost shell; largely determine chemical bonding and reactivity.

Core electrons

Electrons in inner shells; resemble noble gas configurations and are not typically involved in bonding.

Noble gas configuration

Abbreviated electron configuration using a noble gas core to represent inner electrons.

Isoelectronic

Species with the same electron configuration.

Ionization energy (IE1, IE2, …)

Energy required to remove successive electrons from a gaseous atom; IE1 is the first ionization energy.

Electron affinity (EA)

Energy change when adding an electron to a gaseous atom to form an anion; can be negative or positive.

Periodic law

The properties of the elements are periodic functions of their atomic numbers.

Periodic table groups and periods

Groups are vertical columns (1–18); periods are horizontal rows.

Ionic vs covalent bonds

Ionic bonds arise from transfer of electrons; covalent bonds arise from sharing electrons.

Cation vs anion

Cation is a positively charged ion; anion is a negatively charged ion.

Polyatomic ions

Ions composed of more than one atom with an overall charge (e.g., NO3−, SO4^2−).

Oxyanion naming (ate/ite, per-, hypo-)

System for naming oxyanions; ate vs ite indicate oxygen count; per- and hypo- indicate more or fewer oxygens.

Electron configuration exceptions (Cu, Cr)

Deviations due to stability of half-filled or filled d subshells (e.g., Cu: [Ar]4s^1 3d^10; Cr: [Ar]4s^1 3d^5).

Subshell energy order (s < p < d < f)

Within a given principal shell, subshells increase in energy from s to f.

Orbital capacity

Maximum electrons per subshell: s=2, p=6, d=10, f=14.

Orbital vs orbital diagram

Orbital is a region where electrons are likely to be found; orbital diagrams show occupancy with arrows indicating ms.

Transition metals vs inner transition metals

Transition metals: last electron enters d; valence includes ns and (n−1)d. Inner transition metals: last electron enters f; valence includes ns, (n−2)f, and possibly (n−1)d.