Electronic Config. Atoms & Periodic Table (CHEM 201) Chapter 3

Vocabulary flashcards covering fundamental concepts of atomic orbitals, quantum numbers, orbital notation, Aufbau/Pauli/Hund rules, and periodic-table block classifications.

ms (Spin quantum number)

The fourth quantum number for an electron; takes values +1/2 or -1/2 and denotes spin orientation.

Orbital

A region around the nucleus where there is a high probability of finding an electron.

Subshell

A set of orbitals within a shell having the same n and l (e.g., 2p); each subshell has a fixed maximum electron capacity.

Shell

A principal energy level described by the quantum number n; contains one or more subshells.

Degenerate

Orbitals within the same subshell that have the same energy.

Aufbau principle

Electrons fill the lowest-energy orbitals first before occupying higher-energy ones.

Pauli exclusion principle

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

Hund's rule

Electrons in degenerate orbitals fill to maximize the number of unpaired electrons with parallel spins.

Energy diagram

A diagram showing the relative energies of atomic orbitals to visualize ordering and gaps.

Orbital box diagram

A schematic where each orbital is a box and arrows represent electron spins; used to track occupancy and spin.

spdf electronic configuration

Notation that lists electrons by subshells (s, p, d, f) in order of increasing energy, e.g., 1s2 2s2 2p6.

Noble gas shorthand

A compact electron configuration that uses [noble gas] core, followed by the valence subshells (e.g., [Ar] 4s2 3d6).

Maximum electrons in a subshell

s: 2, p: 6, d: 10, f: 14 electrons.

Maximum electrons in a shell

2n^2 electrons; for example, n=1 holds 2, n=2 holds 8, n=3 holds 18, n=4 holds 32.

Periodic table blocks

The table is divided into s-block, p-block, d-block, and f-block corresponding to subshell capacities.

Main group elements

Elements in groups 1A–8A; occupy s- and p-blocks and include representative elements.

Alkali metals

Group 1A elements; highly reactive metals that form bases with water.

Alkaline earth metals

Group 2A elements; form bases and are less soluble in water compared to alkali metals.

Halogens

Group 7A elements; highly reactive nonmetals known as salt-forming elements.

Noble gases

Group 8A elements; very inert due to complete valence shells.

Transition elements

Group B; occupy the d-block and lie between the main group metals and nonmetals.

Chromium and Copper exceptions

Actual configurations differ from the simple Aufbau predictions: Cr prefers [Ar] 4s1 3d5; Cu prefers [Ar] 4s1 3d10 due to stability of half-filled or fully-filled d subshells.

Ground state vs excited state

Ground state: electrons occupy the lowest-energy arrangement. Excited state: electrons promoted to higher-energy orbitals.

Diamagnetic

Materials with all electrons paired; weakly repelled by a magnetic field.

Paramagnetic

Materials with unpaired electrons; attracted to an external magnetic field.

3d orbitals

Five degenerate d orbitals in a given d-subshell: dxy, dyz, dxz, dx^2−y^2, dz^2.

Valence electrons

Electrons in the outermost shell that participate in bonding and chemical reactivity.

Energy ordering within a shell

Within the same principal shell, subshell energies increase with angular momentum: Es < Ep < Ed < Ef; exceptions occur (e.g., 4s fills before 3d).