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 electron configurations.

• 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.

• Periodic law states that the properties of elements recur periodically as a function of their atomic number.

Example: Alkali metals (group 1) all 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 ($m_s$) is a property of the electron and can have values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

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

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

Writing Electron Configurations

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

• Orbital diagram: Uses boxes to represent orbitals and arrows for electrons, showing their spins.

• Short-hand (noble gas) notation: Uses the previous noble gas in brackets to represent core electrons (e.g., [Ne] $3s^2 3p^5$).

• Exceptions: Some d-block elements (e.g., Cu, Ag) have electron configurations that differ from the expected pattern due to stability of filled or half-filled d subshells.

Example: Copper (Cu) has the configuration [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 two 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: For sodium (Na), core electrons are those in the $1s^2 2s^2 2p^6$ orbitals; the valence electron is in $3s^1$.

Properties of Atoms

Periodic Trends

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

• 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: Higher principal quantum number ($n$) means larger orbitals, so atomic size increases down a group.

Magnetic Properties

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

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

Example: Oxygen ($1s^2 2s^2 2p^4$) is paramagnetic due to two unpaired electrons in the $2p$ orbitals.

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

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: For $Na^+$, electron configuration is $1s^2 2s^2 2p^6$ (same as neon).

Periodic Trends Table

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.