Periodicity and Element Properties

The periodic table is organized into groups and periods, with elements sharing similar properties within groups. This page explores the characteristics of different element types:

Metallic Elements:

• Examples: Li, Be, Na, Mg, Al, K, Ca
• Properties: High density, high melting/boiling points, good conductors

Covalent Network Elements:

• Examples: B, C, Si (diamond/graphite)
• Properties: High density, high melting/boiling points, variable conductivity

Monatomic Elements:

• Examples: He, Ne, Ar
• Properties: Low density, low melting/boiling points, non-conductors

Covalent Molecular Elements:

• Examples: H₂, N₂, O₂, F₂, Cl₂, P₄, S₈, fullerenes (C₆₀)
• Properties: Low density, low melting/boiling points, non-conductors

Highlight: Understanding these element categories helps predict their physical and chemical properties, which is crucial for GCSE and higher-level chemistry studies.

Trends in the Periodic Table

The periodic table exhibits several important trends that help explain element properties and behavior:

Covalent Radius: This trend decreases across a period and increases down a group.

Definition: Covalent radius is the distance measured between two nuclei in a shared bond.

Factors influencing covalent radius:

• Across a period: Increased nuclear charge pulls outer electrons closer.
• Down a group: Increased shielding pushes outer electrons away from the nucleus.

Electronegativity: This trend increases across a period and decreases down a group.

Vocabulary: Electronegativity is a measure of an atom's attraction for the shared pair of electrons in a covalent bond.

Factors affecting electronegativity:

• Across a period: Increased nuclear charge enhances attraction to shared electrons.
• Down a group: Increased shielding reduces attraction to shared electrons.

Ionization Energy: This trend generally increases across a period and decreases down a group.

Definition: Ionization energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state.

Factors influencing ionization energy:

• Across a period: Increased nuclear charge makes it harder to remove electrons.
• Down a group: Increased shielding makes it easier to remove outer electrons.

Example: There's a significant jump in ionization energy between the 4th and 5th ionization of silicon because the 5th electron comes from a completed outer shell closer to the nucleus.