Chapter 8: Periodic Properties of the Elements – CHEM 1000 Study Notes

Periodic Properties of the Elements

Introduction

This chapter explores how the arrangement of elements in the periodic table reflects recurring chemical and physical properties. Understanding these trends is essential for predicting element behavior and rationalizing atomic structure.

The Periodic Table

Development and Structure

• Historical Origin: The periodic table was first organized by atomic mass (Mendeleev), but is now arranged by atomic number and electronic structure.

• Periodic Law: Properties of elements recur periodically when arranged by atomic number.

• Application: The periodic table helps determine the electron configuration of each element.

Valence and Core Electrons

Valence Electrons

• Definition: Electrons in the outermost principal energy level; most important for chemical bonding.

• Main Group Elements: Valence electrons are those in the highest principal energy level.

• Transition Elements: Outermost d electrons are also counted as valence electrons.

• Periodic Table Columns: Elements in the same column have the same number of valence electrons.

Core Electrons

• Definition: Electrons in complete principal energy levels and complete d and f sub-levels.

• Example (Silicon): $1s^2 2s^2 2p^6$ are core electrons; $3s^2 3p^2$ are valence electrons.

Orbital Blocks in the Periodic Table

Organization and Electron Configuration

• s-block: 2 columns; single s orbital with 2 electrons.

• p-block: 6 columns; three p orbitals with 2 electrons each.

• d-block: 10 columns; five d orbitals with 2 electrons each.

• f-block: 14 columns; seven f orbitals with 2 electrons each.

Row Number and Principal Quantum Number

• The row number equals the principal quantum number (n) of the highest energy level for main group elements.

• Example (Cl): $1s^2 2s^2 2p^6 3s^2 3p^5$; highest principal level is $n=3$.

Electron Configuration Practice

Example: Germanium (Ge)

• Electron Configuration: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10} 4p^2$

• Valence Electrons: $4s^2 4p^2$ and $3d^{10}$ (for transition elements)

• Core Electrons: All others

Condensed Electron Configuration

• Example (Se): [Ar]$4s^2 3d^{10} 4p^4$

Exceptions to Electron Configurations

• Chromium (Cr): [Ar]$4s^1 3d^5$

• Copper (Cu): [Ar]$4s^1 3d^{10}$

• These exceptions arise due to the close energy levels of 3d and 4s orbitals.

Quantum-Mechanical Model and Periodic Properties

Explanatory Power

• The number of valence electrons determines chemical properties.

• Periodic trends arise because the number of valence electrons is periodic.

Formation of Ions and Predictable Charges

Group 1 and 2 Metal Ions

• Have 1 or 2 electrons more than a noble gas configuration.

• Lose these electrons to form stable cations (e.g., Na+, Mg2+).

Group 16 and 17 Non-Metal Ions

• Have 1 or 2 electrons less than a noble gas configuration.

• Gain electrons to form stable anions (e.g., Cl-, O2-).

Predictable Charges Table

Group	Common Ion
1	+1
2	+2
13	+3
15	-3
16	-2
17	-1

Sizes of Atoms and Ions

Atomic and Ionic Radii

• Atomic radius: Not sharply defined; inferred from distances in compounds.

• Non-bonding atomic radius (van der Waals): Radius of an atom not bonded to another.

• Bonding atomic radius (covalent): Half the distance between two bonded atoms.

• Metals: Half the distance between adjacent atoms in a crystal.

Trends in Atomic Size

• Down a group: Atomic radius increases (higher principal quantum number, larger orbitals).

• Across a period: Atomic radius decreases (increased nuclear charge, same principal quantum number).

Effective Nuclear Charge (Zeff)

Definition and Calculation

• Effective nuclear charge (Zeff): The net positive charge experienced by valence electrons.

• Formula: $Z_{eff} = Z - S$ Where $Z$ is the actual nuclear charge (atomic number), $S$ is the number of core electrons.

• Example (Li): $Z_{eff} = 3 - 2 = 1$

• Example (Be): $Z_{eff} = 4 - 2 = 2$

Shielding

• Core electrons efficiently shield valence electrons from nuclear charge.

• Valence electrons do not efficiently shield each other.

Atomic Radii of d-Block Elements (Transition Metals)

Trends

• Down a group: Atomic radii increase.

• Across a period: Radii remain roughly constant due to increased nuclear charge balanced by increased screening electrons.

• Lanthanide Contraction: Third-row transition elements are similar in size to second-row due to poor shielding by 4f electrons.

Ionic Radius

Cations

• Formed by loss of electrons; nuclear charge remains the same.

• Attraction increases, so ionic radius decreases.

• Example: Li atom (152 pm) vs. Li+ ion (90 pm).

Anions

• Formed by gain of electrons; nuclear charge remains the same.

• Attraction decreases, so ionic radius increases.

• Example: F atom (72 pm) vs. F- ion (119 pm).

Isoelectronic Species

Definition and Trend

• Species with the same number of electrons but different numbers of protons.

• Greater nuclear charge results in a smaller atom or ion.

Species	Number of Protons	Radius (pm)
S2-	16	184
Cl-	17	181
K+	19	133
Ca2+	20	99

Summary Table: Periodic Trends

Property	Down a Group	Across a Period
Atomic Radius	Increases	Decreases
Ionization Energy	Decreases	Increases
Electron Affinity	Generally less negative	Generally more negative
Metallic Character	Increases	Decreases

Key Equations

• Effective Nuclear Charge: $Z_{eff} = Z - S$

• Condensed Electron Configuration: Example: [Ar]$4s^2 3d^{10} 4p^4$ (for Se)

Additional info:

• These notes are based on CHEM 1000, York University, Chapter 8 slides for Fall 2025.

• Recommended problems are provided for further practice (see slide 3).