Chapter 2: Atomic Structure & Interatomic Bonding (Review Flashcards)

Vocabulary-style flashcards covering core concepts from the notes on atomic structure and interatomic bonding.

Atomic number (Z)

Number of protons in the nucleus; equals the number of electrons in a neutral atom.

Isotope

Atoms of the same element with different numbers of neutrons, producing different atomic masses.

Atomic mass unit (amu)

1/12 the mass of carbon-12; 1 amu per atom equals 1 g/mol.

Avogadro's number (N_A)

6.022 × 10^23 particles per mole.

Valence electrons

Electrons in the outermost shell that determine bonding and many properties.

Electron configuration

Arrangement of electrons in shells and subshells (e.g., 1s2 2s2 2p6).

Pauli exclusion principle

No orbital can hold more than two electrons with opposite spins.

Principal quantum number (n)

Shell index (K, L, M, N, …) indicating energy level.

Azimuthal quantum number (l)

Subshell type (s, p, d, f) with values 0 ≤ l ≤ n−1.

Magnetic quantum number (m_l)

Indicates the number and orientation of orbitals within a subshell (values from −l to +l).

Spin quantum number (m_s)

Electron spin: +1/2 or −1/2.

Electronegativity (χ)

Tendency of an atom to attract electrons; Pauling scale ~0.7 to 4.0.

Ionic bonding

Bonding via transfer of electrons creating cations and anions with Coulombic attraction.

Covalent bonding

Bonding via sharing of valence electrons; highly directional.

Metallic bonding

Electron sea model: non-directional bonding leading to close packing, conductivity, and ductility.

Secondary (Van der Waals) bonds

Weaker attractions due to dipole-dipole or induced dipole interactions; always present.

Equilibrium separation (r0)

Interatomic distance at which net energy is minimized and net force is zero.

Attractive, repulsive, and net energies (EA, ER, EN)

EA = attractive energy, ER = repulsive energy, EN = EA + ER; EN minimized at r0.

Bond energy (E0)

Energy required to stretch/break a bond; larger E0 means stronger bond and often higher melting temperature.

Melting temperature (Tm) and bonding

Tm tends to increase with higher bond energy and shorter bond length.

Coulombic attraction in ionic bonding

Main attractive force; proportional to product of ion charges divided by r^2.

Shells and subshells (n and l)

n = energy level (K, L, M, N, …); l = subshell (s, p, d, f). Energy ordering affects occupancy.

Electron states and quantum numbers (ml, ms)

ml selects orbitals within a subshell; ms indicates electron spin orientation.

Periodic table: electropositive vs electronegative elements

Left-side elements tend to lose electrons (form + ions); right-side elements tend to gain electrons (form − ions).

Stability and closed shells

Filled valence shells confer stability; noble gases have stable electron configurations.

Covalent-ionic mixed bonding

Most materials show some ionic/covalent character; percent ionic character rises with electronegativity difference.

Primary vs. Secondary bonds in terms of strength

Primary bonds (ionic, covalent, metallic) are stronger; Secondary (Van der Waals) are weaker.

Examples of ionic compounds

NaCl, MgO, CaF2, CsCl illustrate ionic bonding.

Covalent bonding example

CH4 exemplifies covalent bonding with shared C–H electrons; diamond is a covalent network.

Bonding and material properties

Bond type and bond energy influence melting temperature, thermal expansion, and elastic modulus.

Energy vs. distance concept

Bonding energy curves show a minimum at r0 where the net energy is lowest.

Noble gas electron configuration relevance

Stable configurations resemble closed-shell noble gases and are less reactive.

Energy ordering within a shell

Within a shell, energy increases with the azimuthal quantum number l (and can show overlaps between shells, e.g., 4s vs 3d).