Electron Configuration and Reactivity

Flashcards covering concepts related to electron configuration, orbital diagrams, Hund's rule, valence and core electrons, noble gas configuration, and exceptions for transition metals like chromium and copper, along with ion electron configurations.

P Orbitals

Three atomic orbitals oriented along perpendicular axes that can hold a total of six electrons (two electrons each).

Orbital Diagram

A visual representation of atomic orbitals as dashed lines or boxes, showing all orbitals available for electrons and indicating their spin (ms) and spatial orientation (ml) quantum numbers.

Electron Configuration

A notation describing the distribution of electrons in an atom's orbitals, typically showing the principal quantum number (n) and the angular momentum quantum number (l).

Dash Line Representation

In orbital diagrams, a single dash line represents one s orbital, three dash lines represent the three p orbitals, and five dash lines represent the five d orbitals.

Hund's Rule

States that electrons will occupy each orbital within a subshell singly with parallel spins before any orbital is doubly occupied.

Energy Levels (Orbital Gaps)

The observation that the energy gap between electron orbitals tends to decrease as electrons move further away from the nucleus.

Reactivity (Electron Configuration)

The chemical behavior of an element, which is directly influenced by the distribution, pairing, and location of its electrons, particularly the valence electrons.

Core Electrons

Inner-shell electrons that are completely filled, located closer to the nucleus, stable, and generally do not participate in chemical reactions.

Valence Electrons

Electrons in the highest principal energy shell (valence shell) that determine an atom's chemical properties and are primarily involved in chemical bonding.

Valence Shell

The outermost electron shell of an atom, identified by the highest principal quantum number (n) in its electron configuration, and not always the last orbital filled.

Noble Gas Configuration

A shorthand notation for electron configuration where the symbol of the preceding noble gas in brackets represents the core electrons.

Group Reactivity

The tendency for elements in the same vertical group on the periodic table to exhibit similar chemical reactions due to having the same number of valence electrons.

Chromium and Copper Exceptions

Anomalous electron configurations for these transition metals (and their families) where an electron from the s orbital is promoted to a d orbital to achieve more stable half-filled (d5) or fully-filled (d10) d-subshells.

Half-filled d-orbital

A d-subshell where all five d-orbitals are occupied by a single electron (e.g., d5), resulting in increased stability.

Fully-filled d-orbital

A d-subshell where all five d-orbitals are occupied by two electrons each (e.g., d10), resulting in increased stability.

Electron Promotion

The movement of an electron from a lower energy orbital (like 4s) to a slightly higher energy orbital (like 3d) to achieve a more stable half-filled or fully-filled subshell.

Ions Electron Configuration (Cations)

Derived by first writing the neutral electron configuration, then removing electrons from the valence shell (orbitals with the highest 'n' value) first to account for the positive charge.

Ions Electron Configuration (Anions)

Derived by first writing the neutral electron configuration, then adding electrons to the valence shell to account for the negative charge.

Electron Configuration Importance

A fundamental concept that is crucial for understanding atomic properties like size, ionization energy, electronegativity, and the overall chemical reactivity of elements.

S-D Orbital Energy Difference

The minimal energy difference between s and d orbitals (especially for higher principal quantum numbers) that justifies the promotion of electrons between them to achieve greater stability.