Topic 1: Atomic Structure and Periodicity

Vocabulary flashcards covering key terms from Topic 1: Atomic Structure and Periodicity in CHM012.

Electromagnetic radiation (EMR)

Energy that travels through space; visible light is one type of EMR.

Wavelength (λ)

Distance between two consecutive peaks or troughs of a wave; measured in nanometers (nm).

Frequency (ν)

Number of wave cycles per second; unit is hertz (s⁻¹).

Speed of light (c)

Constant 2.9979 × 10^8 m/s; c = λν.

Planck’s constant (h)

6.626×10⁻³⁴ J·s; relates energy to frequency via E = hν.

Photon

Quantum of electromagnetic radiation; energy E = hν.

Quantum (of energy)

Energy transfer occurs in whole-number quanta, giving EMR particle-like properties.

Energy quantization

Energy can be gained or lost only in whole-number multiples of hv.

Bohr model

Electrons occupy allowed energy levels; ΔE = hv; works best for hydrogen.

Allowed energy level

Specific energy values electrons may occupy in an atom.

Degenerate orbitals

Orbitals that have the same energy (in hydrogen, orbitals with same n).

Orbital

Region in space described by quantum numbers where there is a high probability of finding an electron.

Principal quantum number (n)

Descriptors of energy level; integers n ≥ 1.

Angular momentum quantum number (l)

Determines orbital shape; 0 ≤ l ≤ n−1; s (0), p (1), d (2), f (3).

Magnetic quantum number (ml)

Describes orbital orientation; −l ≤ ml ≤ l.

Spin quantum number (ms)

Electron spin; values +½ or −½.

Pauli Exclusion Principle

No two electrons in an atom can have identical four quantum numbers.

Hund’s Rule

For degenerate orbitals, maximize unpaired electrons with the same spin before pairing.

Electron configuration

Arrangement of electrons in orbitals; ground state is lowest energy.

Valence electrons

Electrons in the outermost shell; determine chemical properties.

Core electrons

Electrons in inner shells that are not valence.

Noble gas (condensed) configuration

Shortened electron configuration using [noble gas] to represent filled inner shells.

Orbital diagrams

Diagram of orbitals as boxes with arrows indicating electrons and their spins.

s, p, d, f orbitals

Types of orbitals: s (l=0, sphere), p (l=1, two lobes), d (l=2, four lobes plus donut), f (l=3, complex).

Aufbau principle

Electrons fill orbitals in order of increasing energy.

Electron configuration anomalies (Cr, Cu)

Some elements show unexpected configurations due to close 4s/3d energies (e.g., Cr: [Ar] 4s¹ 3d⁵).

Periodic properties (periodicity)

Repeating trends in properties with atomic number (radius, ionization energy, electron affinity).

Effective nuclear charge (Zeff)

Zeff = Z − S; Zeff increases across a period and decreases down a group.

Atomic radius

Half the distance between nuclei in a bonded pair; decreases across a period, increases down a group.

Ionic radius

Size of ions; depends on nuclear charge and electron count; cations smaller, anions larger; isoelectronic series shows radius decreasing with increasing nuclear charge.

Isoelectronic series

Ions with the same number of electrons; radius decreases as nuclear charge increases.

Ionization energy (IE)

Energy required to remove an electron from a gaseous atom/ion; IE1 is for the first electron; generally increases across a period and decreases down a group.

Electron affinity (EA)

Energy change when adding an electron to a gaseous atom; typically exothermic (negative), with notable exceptions.

Hydrogen spectrum vs continuous spectrum

Continuous spectrum contains all visible wavelengths; hydrogen line spectrum shows discrete lines.