Electron Configuration

Introduction to Electron Configuration

Electron configuration describes the arrangement of electrons in an atom's orbitals. Understanding electron configuration is essential for predicting chemical properties, reactivity, and the placement of elements in the periodic table.

• Periodic Law: When elements are arranged by increasing atomic number, elements in the same group share similar chemical properties.

• Orbitals: Regions in an atom where electrons are likely to be found. Types include s, p, d, and f orbitals.

• Electron Arrangement: The order in which electrons fill atomic orbitals is determined by their energy levels.

Rules for Electron Configuration

• Aufbau Principle: Electrons fill orbitals starting from the lowest energy level to the highest.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers; each orbital can hold a maximum of two electrons with opposite spins.

• Hund's Rule: Electrons occupy degenerate (equal energy) orbitals singly before pairing up.

Notation and Examples

• Standard Notation: Lists each occupied subshell and the number of electrons in it. For example, for fluorine (Z=9):

• Condensed (Noble Gas) Notation: Uses the symbol of the previous noble gas in brackets to represent core electrons. For example, aluminum (Z=13):

• Example: Zinc (Z=30):

Blocks of the Periodic Table

• s-block: Groups 1A-2A (alkali and alkaline earth metals)

• p-block: Groups 3A-8A (main group elements)

• d-block: Transition metals (groups 3B-12B)

• f-block: Lanthanides and actinides (bottom two rows)

Exceptions to Electron Configuration

Certain elements have electron configurations that differ from the expected pattern due to increased stability of half-filled or fully filled d and f subshells.

• Common Exceptions: Chromium (Cr), Molybdenum (Mo), Copper (Cu), Silver (Ag), Gold (Au), and others.

• Example: Silver (Ag, Z=47): Expected: Actual:

• Reason: Half-filled and fully filled subshells are more stable due to exchange energy and symmetry.

Electron Configurations for Ions

• Cations (positive ions): Electrons are removed first from the highest principal quantum number (n) orbital.

• Anions (negative ions): Electrons are added to the next available orbital following the Aufbau principle.

• Example: Nitride ion (): (same as neon, a noble gas configuration)

Paramagnetic vs. Diamagnetic

• Paramagnetic: Atoms or ions with at least one unpaired electron; attracted to magnetic fields.

• Diamagnetic: Atoms or ions with all electrons paired; weakly repelled by magnetic fields.

• Example: Vanadium (V, Z=23): (has unpaired electrons, so it is paramagnetic)

Quantum Numbers

Introduction to Quantum Numbers

Quantum numbers describe the properties of atomic orbitals and the properties of electrons in those orbitals. Each electron in an atom is described by a unique set of four quantum numbers.

Assigning Quantum Numbers

• Step 1: Write the electron configuration for the atom or ion.

• Step 2: Identify the orbital for the electron of interest.

• Step 3: Assign values for n, l, ml, and ms based on the orbital and electron spin.

• Example: Oxygen (O, Z=8): Electron configuration: For a 2p electron: n = 2, l = 1, ml = -1, 0, or +1, ms = +1/2 or -1/2

Application: Determining Electron Location

• Each electron in an atom is uniquely identified by its set of quantum numbers.

• Quantum numbers determine the size, shape, and orientation of orbitals, as well as the spin of electrons.

Summary Table: Electron Configuration and Quantum Numbers

Additional info: Some content was inferred and expanded for clarity, including the explanation of quantum numbers, the periodic table blocks, and the summary tables, to ensure the notes are self-contained and academically robust.