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 into groups (columns) and periods (rows), with elements classified as main-group or transition elements based on their position and electron configuration.

• Main-group elements are found in groups 1, 2, and 13–18. They include the s-block and p-block elements.

• Transition elements are located in groups 3–12 and correspond to the d-block of the periodic table.

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

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

Electron Configuration of Elements

Spin Quantum Number and Electron Properties

Electrons possess a property called spin, described by the spin quantum number (ms).

• The spin quantum number can have values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

• This property is intrinsic to electrons and affects their arrangement in orbitals.

Pauli Exclusion Principle

The Pauli Exclusion Principle states that 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.

Writing Electron Configurations

• Long-hand electron configuration: List all occupied orbitals using spdf notation (e.g., $1s^2 2s^2 2p^6$).

• Orbital diagram: Use boxes to represent orbitals and arrows for electrons, following Hund’s rule.

• Hund’s Rule: Electrons fill degenerate orbitals singly before pairing.

• Short-hand (noble gas) notation: Use the previous noble gas in brackets, then add remaining configuration (e.g., $[Ne] 3s^2 3p^5$).

Example: Chlorine: $1s^2 2s^2 2p^6 3s^2 3p^5$ or $[Ne] 3s^2 3p^5$

Exceptions in d-block Elements

• Some d-block elements (e.g., Copper (Cu) and Silver (Ag)) have electron configurations that deviate from the expected pattern to achieve greater stability.

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

Blocks of the Periodic Table

• s-block: Groups 1 and 2, plus Helium

• p-block: Groups 13–18

• d-block: Groups 3–12 (transition metals)

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

Core and Valence Electrons

• Core electrons: Electrons in inner shells, not involved in bonding.

• Valence electrons: Electrons in the outermost shell, responsible for chemical properties.

Example: Sodium ($Na$): $1s^2 2s^2 2p^6 3s^1$ – 1 valence electron ($3s^1$), 10 core electrons.

Properties of Atoms

Periodic Trends

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

• Atomic size (radius): 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 (stronger attraction for electrons).

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

Explanation of Trends

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

• Orbital size: As principal quantum number increases, orbitals become larger, affecting atomic size and other properties.

Example: Fluorine has a higher ionization energy than sodium due to higher $Z_{eff}$ and smaller atomic radius.

Magnetic Properties: Diamagnetic vs. Paramagnetic

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

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

Example: Oxygen ($O_2$) is paramagnetic due to unpaired electrons.

Predicting Charges of Main Group Elements

• Main-group elements tend to form ions with charges that achieve noble gas configurations.

• Group 1: $+1$; Group 2: $+2$; Group 17: $-1$; Group 16: $-2$

Example: Magnesium forms $Mg^{2+}$ by losing two electrons.

Electron Configuration for Ions

• Remove electrons from the highest energy orbital first when forming cations.

• Add electrons to the lowest available orbital when forming anions.

Example: $Cl^-$: $1s^2 2s^2 2p^6 3s^2 3p^6$ ($[Ar]$ configuration)

Summary Table: Periodic Trends

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

Additional info: Table summarizes the main periodic trends for quick reference.