Electron Configuration and Periodic Properties: Study Notes

Electron Configuration and the Periodic Table

Introduction

The position of an element in the periodic table determines its electron configuration and influences its chemical and physical properties. Elements within the same group (vertical column) tend to have similar properties due to similar valence electron arrangements.

• Group: Vertical column in the periodic table; elements in the same group have similar chemical properties.

• Period: Horizontal row in the periodic table; elements in the same period have the same number of electron shells.

• Example: Elements such as Cl and Br (Group 17) have similar chemical properties.

Classification of Elements

Types of Elements

Elements are classified based on the type of subshell containing the outermost electrons:

• Main group elements (representative elements): Groups 1, 2, and 13-18.

• Transition elements: Groups 3-12; electrons are added to d subshells.

• Inner transition elements: Lanthanides and actinides; electrons are added to f subshells.

Electron configuration notation: $n s^2 n p^6$ (for main group elements)

Electron Configurations of Groups

Group 1 and Group 2

Alkali metals (Group 1) and alkaline earth metals (Group 2) have a single valence electron in the s orbital, leading to similar chemical behavior within each group.

• Alkali metals: $[\text{noble gas}] n s^1$

• Alkaline earth metals: $[\text{noble gas}] n s^2$

• Example: Na: $[\text{Ne}] 3s^1$

Group 17 (Halogens)

• Halogens have seven valence electrons: $[\text{noble gas}] n s^2 n p^5$

• Example: Cl: $[\text{Ne}] 3s^2 3p^5$

Exceptions

• Some groups contain both metals and nonmetals (e.g., Group 14: C is a nonmetal, Sn and Pb are metals).

Effective Nuclear Charge ($Z_{eff}$)

Definition and Calculation

Effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. It is less than the actual nuclear charge due to shielding by other electrons.

• $Z_{eff} = Z - S$

• $Z$ = number of protons

• $S$ = shielding constant (approximate number of inner electrons)

• Example: For hydrogen, $Z_{eff} = 1$; for helium, $Z_{eff} < 2$ due to electron-electron repulsion.

Periodic Trends

Atomic Radius

Atomic radius is the distance between the nucleus of an atom and its valence shell. It can be defined in several ways:

• Van der Waals radius: Half the distance between adjacent, non-bonded atoms.

• Covalent radius: Half the distance between adjacent, identical nuclei in a molecule.

Trends:

• Atomic radius decreases from left to right across a period (due to increasing $Z_{eff}$).

• Atomic radius increases from top to bottom within a group (due to increasing principal quantum number).

Ionization Energy (IE)

Definition and Trends

Ionization energy is the minimum energy required to remove an electron from an atom in the gas phase.

• First ionization energy: $\text{X}(g) \rightarrow \text{X}^+(g) + e^-$

• IE increases across a period (left to right).

• IE decreases down a group.

Exceptions:

• Between Groups 2 and 13, and Groups 15 and 16, due to subshell stability and electron repulsion.

Electron Affinity (EA)

Definition and Trends

Electron affinity is the energy released when an atom in the gas phase accepts an electron.

• EA is generally positive (energy released).

• EA increases (becomes more negative) across a period.

• EA decreases down a group.

Example: $\text{Cl}(g) + e^- \rightarrow \text{Cl}^-(g)$, $EA = -349 \text{ kJ/mol}$

Electron Configurations of Ions

Writing Electron Configurations

To write the electron configuration of an ion, start with the neutral atom and add electrons for anions or remove electrons for cations.

• Example: $\text{Cl}^-: [\text{Ne}] 3s^2 3p^6$

• Example: $\text{Na}^+: [\text{Ne}]$

Isoelectronic Species

Definition and Examples

Isoelectronic species are atoms or ions with identical electron configurations.

• Example: $\text{Na}^+$ and $\text{Ne}$ are isoelectronic: $[\text{Ne}]$

• Example: $\text{Cl}^-$ and $\text{Ar}$ are isoelectronic: $[\text{Ne}] 3s^2 3p^6$

Ionic Radius

Definition and Trends

Ionic radius changes when an atom gains or loses electrons:

• Anions (gain electrons): radius increases due to increased electron-electron repulsion.

• Cations (lose electrons): radius decreases due to reduced electron-electron repulsion.

• Example: $\text{Na}: 186 \text{ pm}$, $\text{Na}^+: 102 \text{ pm}$

Orbital Energies

Single-Electron and Multi-Electron Systems

Energy of orbitals depends on the principal quantum number ($n$) and angular momentum quantum number ($l$).

• For hydrogen-like atoms, energy increases with $n$.

• For multi-electron atoms, energy also depends on $l$; $s < p < d < f$ for a given $n$.

• Relative energy: $4s < 3d$

Pauli Exclusion Principle

Definition

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

• Symbolic notation: $n l^{\text{number of electrons}}$

Aufbau Principle

Definition and Application

Electrons fill orbitals in order of increasing energy, starting with the lowest energy orbital available.

• Order: $1s \rightarrow 2s \rightarrow 2p \rightarrow 3s \rightarrow 3p \rightarrow 4s \rightarrow 3d$

Hund's Rule

Definition

The most stable arrangement of electrons in orbitals of equal energy (degenerate orbitals) is the one with the maximum number of unpaired electrons with parallel spins.

• Electrons occupy empty orbitals singly before pairing up.

• Maximizes electron-electron repulsion, increasing stability.

Electron Configuration Examples and Anomalies

Examples

• Titanium (Ti, Z=22): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^2$

• Chromium (Cr, Z=24): $[\text{Ar}] 4s^1 3d^5$ (half-filled d orbital is more stable)

• Copper (Cu, Z=29): $[\text{Ar}] 4s^1 3d^{10}$ (completely filled d orbital is more stable)

HTML Table: Periodic Trends Summary

Trend	Across a Period (Left to Right)	Down a Group (Top to Bottom)
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	Increases (more negative)	Decreases
Effective Nuclear Charge ($Z_{eff}$)	Increases	Decreases slightly

Additional Info

• Transition metals and inner transition metals have unique electron configurations due to d and f subshell filling.

• Lanthanides and actinides are characterized by incomplete f subshells.

• Isoelectronic species are important for understanding chemical reactivity and ionic sizes.