Electronic Structure and Quantum Numbers - Vocabulary Flashcards

These vocabulary flashcards cover key terms and concepts from the lecture notes on electronic structure, subshells, electron configurations, LEDS, and quantum numbers.

Shell (Energy level)

An energy level of electrons around the nucleus, labeled by the principal quantum number n.

Subshell

Sublevels within a shell; indicate the shape of orbitals and are labeled s, p, d, or f.

Orbital

A region within a subshell where an electron is likely to be found; each orbital holds up to 2 electrons with opposite spins.

Principal Quantum Number (n)

Describes the energy level of an electron; larger n means farther from the nucleus.

Azimuthal Quantum Number (l)

Describes orbital shape; values range from 0 to n−1 and map to s (0), p (1), d (2), f (3).

s orbital

Orbital with l = 0; describes a spherical shape; there is 1 orbital in the subshell.

p orbital

Orbital with l = 1; describes a dumbbell-shaped region; there are 3 orbitals in the subshell.

d orbital

Orbital with l = 2; describes a more complex shape; there are 5 orbitals in the subshell.

f orbital

Orbital with l = 3; complex shape; there are 7 orbitals in the subshell.

Max electrons in s subshell

2 electrons.

Max electrons in p subshell

6 electrons.

Max electrons in d subshell

10 electrons.

Max electrons in f subshell

14 electrons.

Orbitals per subshell

S subshell has 1 orbital; P subshell has 3 orbitals; D subshell has 5 orbitals; F subshell has 7 orbitals.

Electron Configuration

The arrangement of electrons in an atom’s energy levels and sublevels.

Full Notation (Electron Configuration)

Writing the complete configuration, e.g., Na: 1s2 2s2 2p6 3s1.

Noble Gas Notation

Shorthand electron configuration using the nearest previous noble gas, e.g., Cl: [Ne] 3s2 3p5.

Valence Electrons

Electrons in the outermost energy level that participate in bonding; e.g., Na has 1 valence electron.

Valence electron count (example)

Na has 1 valence electron; Cl has 7 valence electrons.

Aufbau Principle

Electrons fill from the lowest energy level upward in order of increasing energy.

Pauli Exclusion Principle

Each orbital can hold at most 2 electrons with opposite spins.

Hund’s Rule

Electrons fill degenerate orbitals singly before pairing to minimize repulsion.

Paramagnetism

Attracted to a magnetic field due to unpaired electrons.

Diamagnetism

Slightly repelled by a magnetic field due to all electrons being paired.

Octet Rule

Atoms tend to have 8 electrons in their outer shell (except H and He for the duet rule).

Duet Rule

H and He prefers 2 valence electrons instead of 8.

Lewis Electron Dot Structure (LEDS)

A diagram showing valence electrons as dots around the element symbol.

LEDS for Neutral Atoms

Place one dot on each side before pairing until all valence electrons are represented.

LEDS for Ions

Enclose the atom in brackets with a charge outside; add dots to complete octet as needed.

Ionic Bond

Bond formed by transfer of electrons between atoms.

Covalent Bond

Bond formed by sharing of electrons between atoms.

Central Atom (in LEDS/covalent molecules)

Typically the least electronegative element in a molecule.

Electronegativity

Tendency of an atom to attract electrons toward itself; generally higher to the right and upper parts of the periodic table.

Periodic Table Blocks

s-block, p-block, d-block, and f-block correspond to where valence electrons reside.

Metals vs Nonmetals

Left side of the periodic table generally metals; right side nonmetals; metalloids along the border.

Noble Gases

Group 18 elements with full outer electron shells; generally unreactive (He is an exception with 2 valence electrons).

[Ar] notation

An example of noble gas shorthand, e.g., Ca: [Ar] 4s2.

Group 1 Alkali Metals

Elements with 1 valence electron; highly reactive.

Magnetic Quantum Number (ml)

Orientation of the orbital in space; number of orbitals in a subshell; ranges from −l to +l; 2l+1 orbitals.

Spin Quantum Number (ms)

Electron spin state; either +1/2 or −1/2.

Atomic Number

The number of protons in the nucleus; defines the identity of the element.

Orbital Occupancy Rule (2l+1)

The number of orbitals in a subshell is 2l + 1 and equals the range of ml values.

S-block vs P-block electrons

S-block electrons fill s subshells; P-block electrons fill p subshells; their blocks define portions of the periodic table.

Argon example for shorthand configuration

[Ar] 4s2 is a common shorthand for calcium's electron configuration.

Electron Shells and Subshells

Shells are energy levels; subshells (s, p, d, f) are subdivisions within shells that determine electron distribution.

Shape of s-sublevel

Spherical shape corresponding to l = 0.

Shape of p-sublevel

Dumbbell-shaped orbitals corresponding to l = 1.

Shape of d-sublevel

Complex shapes corresponding to l = 2.