Chapter 3 Study Notes: Electron Configuration, Orbital Diagrams, and Periodic Properties

Electron Configuration and Orbital Diagrams

Ground State Electron Configuration

The ground state electron configuration of an atom describes the arrangement of electrons in its atomic orbitals, following the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Aufbau Principle: Electrons fill orbitals starting with the lowest energy level first.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

Example: The ground state electron configuration of Phosphorus (P) is:

• $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^3$

Orbital Diagrams

An orbital diagram uses arrows to represent electrons in each orbital, showing their spin and arrangement.

• Each box represents an orbital; arrows indicate electrons and their spins.

• Example for Phosphorus (P):

1s: ↑↓ 2s: ↑↓ 2p: ↑↓ ↑↓ ↑↓ 3s: ↑↓ 3p: ↑ ↑ ↑

Abbreviated Electron Configuration

Using Noble Gas Notation

The abbreviated electron configuration uses the previous noble gas to simplify notation.

• Example: Calcium (Ca) — [Ar] $4s^2$

• Example: Titanium (Ti) — [Ar] $4s^2 3d^2$

Periodic Trends and Properties

Paramagnetic vs. Diamagnetic

Paramagnetic atoms or ions have unpaired electrons and are attracted to magnetic fields. Diamagnetic species have all electrons paired and are weakly repelled by magnetic fields.

• Example: Al ($1s^2 2s^2 2p^6 3s^2 3p^1$) is paramagnetic (one unpaired electron).

• Cl- ($1s^2 2s^2 2p^6 3s^2 3p^6$) is diamagnetic (all electrons paired).

Effective Nuclear Charge ($Z_{eff}$)

Effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom.

• Calculated as: $Z_{eff} = Z - S$

• Where $Z$ is the atomic number and $S$ is the number of shielding electrons.

• Example: For a valence electron in Na ($Z=11$), $S=10$, so $Z_{eff}=1$.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion.

• First ionization energy increases across a period and decreases down a group.

• Order for Li, Na, K: $Li > Na > K$

Atomic Radius

Atomic radius generally decreases across a period and increases down a group.

• Order for O2-, F-, N3-: $N^{3-} > O^{2-} > F^{-}$

Electron Configuration for Ions

Transition Metal Ions

When transition metals form ions, electrons are removed first from the highest principal quantum number (usually the s orbital).

• Example: Zn2+ — [Ar] $3d^{10}$

Summary Table: Ionization Energies

The following table compares the first, second, and third ionization energies for Na, Mg, and Al.

Element	1st IE (kJ/mol)	2nd IE (kJ/mol)	3rd IE (kJ/mol)
Na	495.8	4562	6912
Mg	737.7	1450.7	7732.7
Al	577.5	1816.7	2744.8

Application: Large jumps in ionization energy indicate removal of a core electron.

Practice Problems

• Write ground state electron configurations for given elements and ions.

• Determine paramagnetic or diamagnetic nature.

• Rank elements by ionization energy and atomic radius.

• Calculate effective nuclear charge for specified atoms.

Additional info: These notes cover key concepts from Chapter 3: Molecules, Compounds & Chemical Equations, with emphasis on electron configuration, periodic trends, and atomic properties relevant for General Chemistry students.