Chapter 3: Periodic Properties of the Elements – Study Guide

Periodic Law and the Periodic Table

Main-Group and Transition Elements

The periodic table is organized to reflect recurring chemical properties. Elements are classified as main-group elements (groups 1, 2, and 13–18) and transition elements (groups 3–12).

• Main-group elements: Located in the s- and p-blocks; include alkali metals, alkaline earth metals, and groups 13–18.

• Transition elements: Located in the d-block; include groups 3–12.

• Similar properties: Elements in the same group (vertical column) have similar chemical and physical properties due to similar valence electron configurations.

Example: All alkali metals (group 1) react vigorously with water and form +1 ions.

Electron Configuration of Elements

Quantum Numbers and Electron Properties

Electrons in atoms are described by four quantum numbers. The spin quantum number ($m_s$) describes the intrinsic spin of the electron and can have values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

• Spin quantum number ($m_s$): Indicates the two possible spin states of an electron.

Pauli Exclusion Principle

The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers. This limits each orbital to two electrons with opposite spins.

• Application: When writing electron configurations, each orbital can hold a maximum of two electrons with opposite spins.

Writing Electron Configurations

Electron configurations describe the arrangement of electrons in an atom. The notation uses the letters s, p, d, and f to indicate subshells.

• Long-hand configuration: Lists all occupied subshells (e.g., $1s^2\ 2s^2\ 2p^6\ 3s^2\ 3p^6\ 4s^2$).

• Orbital diagram: Uses boxes to represent orbitals and arrows for electrons, following Hund’s Rule (electrons fill degenerate orbitals singly before pairing).

• Short-hand (noble gas) notation: Uses the previous noble gas in brackets to represent core electrons (e.g., $[Ne]\ 3s^2\ 3p^6$).

Hund’s Rule: Electrons occupy orbitals singly with parallel spins before pairing.

Exceptions in d-block Elements

Some d-block elements (such as copper and silver) have electron configurations that deviate from the expected pattern to achieve greater stability.

• Example: Copper: $[Ar]\ 4s^1\ 3d^{10}$ instead of $[Ar]\ 4s^2\ 3d^9$.

• Reason: A filled or half-filled d subshell is energetically favorable.

Periodic Table Blocks

• s-block: Groups 1 and 2, plus helium.

• p-block: Groups 13–18.

• d-block: Transition metals, groups 3–12.

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

Core and Valence Electrons

Core electrons are those in filled inner shells; valence electrons are in the outermost shell and participate in chemical bonding.

• Identification: Valence electrons are those beyond the noble gas core in the electron configuration.

Properties of Atoms

Periodic Trends

Several atomic properties change predictably across periods and groups due to electron configuration and effective nuclear charge.

• Atomic size: Decreases across a period, increases down a group.

• Ionization energy: Increases across a period, decreases down a group.

• Ionic radii: Cations are smaller than their parent atoms; anions are larger.

• Electron affinity: Generally becomes more negative across a period.

• Metallic character: Increases down a group, decreases across a period.

Explanation of Trends

Trends are explained by effective nuclear charge ($Z_{eff}$) and orbital size.

• Effective nuclear charge ($Z_{eff}$): The net positive charge experienced by valence electrons. Increases across a period, causing electrons to be held more tightly.

• Orbital size: Larger orbitals (higher principal quantum number) mean electrons are farther from the nucleus and less tightly held.

$Z_{eff} = Z - S$

Where $Z$ is the atomic number and $S$ is the shielding constant.

Magnetic Properties

Atoms can be classified based on their magnetic properties:

• Diamagnetic: All electrons are paired; not attracted to a magnetic field.

• Paramagnetic: Contains unpaired electrons; attracted to a magnetic field.

Predicting Charges of Main Group Elements

Main group elements tend to form ions with charges that result in a noble gas electron configuration.

• Group 1: Forms $+1$ ions.

• Group 2: Forms $+2$ ions.

• Group 17: Forms $-1$ ions.

• Group 16: Forms $-2$ ions.

Electron Configuration for Ions

When atoms form ions, electrons are added or removed from the outermost shell.

• Cations: Electrons are removed from the highest energy orbital.

• Anions: Electrons are added to the lowest available energy orbital.

Example: Sodium ion ($Na^+$): $1s^2\ 2s^2\ 2p^6$ (loses one electron from $3s$).

Summary Table: Periodic Trends

Property	Across Period (Left to Right)	Down Group (Top to Bottom)
Atomic Size	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	Becomes more negative	Becomes less negative
Metallic Character	Decreases	Increases

Additional info: The above table summarizes the main periodic trends for atomic properties.