General Chemistry I: Chapter 6 - Electronic structure and periodic properties of elements

A comprehensive set of VOCABULARY flashcards covering electromagnetic energy, quantum theory, atomic orbitals, quantum numbers, electron configurations, and periodic properties from the notes.

Electromagnetic radiation

Energy that propagates as waves with oscillating electric and magnetic fields; includes visible light and other wavelengths; travels through vacuum at the speed of light.

Speed of light (c)

The constant speed at which electromagnetic waves travel in vacuum, c ≈ 2.998 × 10^8 m/s; related to wavelength and frequency by c = λν.

Wavelength (λ)

Distance between consecutive peaks or troughs in a wave; inversely related to frequency.

Frequency (ν)

Number of wave cycles that pass a point per unit time; measured in hertz (Hz).

Planck’s constant (h)

6.63 × 10^-34 J·s; relates energy and frequency via E = hν; foundational to quantum theory.

Photon

A quantum of light with energy E = hν = hc/λ; exhibits particle-like properties.

Photoelectric effect

Emission of electrons from a material when illuminated with light above a threshold frequency; supports the particle nature of light.

Threshold frequency

Minimum frequency of light required to eject electrons in the photoelectric effect.

Planck’s quantum theory

Energy is emitted or absorbed in discrete units called quanta; E = hν.

Wave-particle duality

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

de Broglie wavelength

Matter waves; λ = h/p, where p is momentum; shows particles have wave-like properties.

Schrödinger equation

Wave equation that describes the quantum state of a system; solutions yield the wavefunction and electron density.

Quantum numbers

Set of numbers (n, l, mℓ, ms) that describe the state of electrons in atoms.

Principal quantum number (n)

Shell or energy level; n = 1, 2, 3, …; relates to distance from the nucleus.

Angular momentum quantum number (l)

Subshell type within a given n (0 = s, 1 = p, 2 = d, 3 = f); determines orbital shape; l ranges from 0 to n−1.

Magnetic quantum number (mℓ)

Orientation of the orbital in space; for a given l, mℓ = −l,…,0,…,+l.

Spin quantum number (ms)

Electron spin; values +1/2 or −1/2; two electrons in the same orbital must have opposite spins.

s orbital

l = 0; spherical orbital; one orbital per energy level; holds up to 2 electrons.

p orbital

l = 1; three degenerate orbitals (px, py, pz); ml = −1, 0, 1; holds up to 6 electrons.

d orbital

l = 2; five orbitals; ml = −2, −1, 0, 1, 2; holds up to 10 electrons.

f orbital

l = 3; seven orbitals; ml = −3, −2, −1, 0, 1, 2, 3; holds up to 14 electrons.

Aufbau principle

Electrons fill the lowest-energy subshells first in building up the electron configuration.

Pauli exclusion principle

No two electrons in an atom can have the same four quantum numbers; a maximum of two electrons per orbital with opposite spins.

Hund’s rule

Electrons fill degenerate orbitals singly with parallel spins before pairing.

Valence electrons

Electrons in the outermost shell that determine chemical properties and reactivity.

Core electrons

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

Isoelectronic

Species that have the same electron configuration; size is determined by nuclear charge within an isoelectronic series.

Noble gas core notation

Abbreviated electron configuration using [noble gas] to represent inner electrons, e.g., [Ne]3s2 3p4.

Zeff (effective nuclear charge)

The net positive charge felt by a valence electron; Zeff ≈ Z − s, where s is shielding; increases across a period, decreases down a group.

Covalent radius

Half the distance between nuclei in a covalent bond of identical atoms; increases down a group, decreases across a period.

Ionic radius

Radius of an ion; cations are smaller than their parent atoms, anions larger.

Isoelectronic trend in radius

In isoelectronic series, atoms/ions with more protons (higher Z) are smaller.

Ionization energy (IE1)

Energy required to remove the first electron from a gaseous atom; positive and generally decreases down a group, increases across a period.

Second ionization energy (IE2)

Energy required to remove the second electron after the first has been removed; typically much larger than IE1.

Electron affinity (EA)

Energy change when a gaseous atom gains an electron to form an anion; can be exothermic (negative EA) or endothermic (positive EA), trends vary among groups.

Rydberg constant (R∞)

1.097 × 10^7 m−1; used in hydrogen emission/ absorption calculations; relates to energy levels in hydrogen.

Balmer series

Hydrogen emission lines in the visible region corresponding to transitions to n1 = 2.

Lyman series

Hydrogen emission lines in the ultraviolet region corresponding to transitions to n1 = 1.

Hydrogen energy level formula (E_n)

E_n = −RH / n^2; energy of the nth level in the hydrogen atom.

Bohr model basics

Electrons in hydrogen occupy quantized energy levels; energy differences produce photons; En = −RH(1/n^2).

Line spectra vs continuous spectra

Line spectra show discrete wavelengths from transitions between levels; continuous spectra arise from broad ranges of wavelengths.

Hydrogen energy transition formula (ΔE)

ΔE = RH(1/n1^2 − 1/n2^2) for hydrogenic transitions; determines photon energy and wavelength.

Hydrogen vs multi-electron atoms

Schrödinger equation solvable exactly only for H; multi-electron atoms require approximation.