Lecture 3 Inorganic Chemistry CHE 211

CHE 221: Inorganic Chemistry

  • Lecture: Basic Science Lecture 3: Periodic Table of Elements

  • Date: 22/10/2024

  • Instructor: Dr. TaReK El-DaBea

Periodic Table of Elements

  • Overview:

    • Organized in rows and columns based on atomic numbers and properties.

  • Elements and Their Properties:

    • Columns represent groups with similar properties.

    • Rows represent periods with increasing atomic numbers.

  • Important Elements:

    • Hydrogen (H): Atomic Number 1, Symbol H, Atomic Weight 1.00794.

    • Helium (He): Atomic Number 2, Symbol He, Atomic Weight 4.002602.

    • Includes names, atomic numbers, weights, symbols for all elements up to Radon and beyond.

Periodic Trends in Properties

Variation in Atomic Radii

  1. Variation in a Period:

    • General decrease in atomic radius across a period for s and p-block elements due to increasing effective nuclear charge ( [Z_{eff}]).

    • Example: Li > Be > B > C > N > O > F > Ne.

    • For d-block elements, atomic radii initially decrease, remain constant, and then increase.

  2. Variation in a Group:

    • Atomic radii increase down a group due to the addition of electron shells.

    • Exceptions:

      • Gallium (Ga) has a radius similar to Aluminum (Al) despite expectations.

      • Thallium (Tl) is similar in radius to Indium (In) contrary to predictions. Causal factors include electronic configurations and shielding effects.

Ionic Radius

  1. Positive Ions:

    • Formed when electrons are removed; result in a smaller ionic size.

    • Example: A+ < A.

  2. Negative Ions:

    • Formed when electrons are added; result in a larger ionic size.

    • Example: A- > A.

  3. Factors Affecting Ionic Radius:

    • Shell number increases leads to an increase in ionic radius.

    • Cation charge increase results in a decrease in ionic radius.

    • Anion charge increase results in a larger ionic radius.

Ionization Energy (IE)

  1. Definition:

    • Energy required to remove an electron from an isolated gaseous atom.

    • Measured in eV atom-1 or kJ mol-1.

  2. Trends:

    • Ionization energy generally increases across a period due to increasing (Z_{eff}).

    • Exceptions include fully-filled and half-filled electronic configurations possessing higher IEs.

  3. Factors:

    • Increases with (Z_{eff}) and decreases when the atomic radius increases.

    • Higher for half-filled and fully-filled configurations.

Electron Affinity

  1. Definition:

    • Energy released when an electron is added to a neutral atom in its gaseous state.

  2. Trends:

    • Electron affinity generally increases across a period; decreases down a group.

    • Notable exceptions exist for p-block elements between periods.

  3. Factors:

    • Increases with greater (Z_{eff}) and decreases with increases in atomic radius.

Electronegativity

  1. Definition:

    • Tendency of an atom to attract electrons towards itself in a compound.

  2. Trends:

    • Increases from left to right across a period and generally decreases down a group.

  3. Noteworthy Points:

    • Fluorine is the most electronegative element (4.0), Cesium the least (0.7).

Hydration Energy

  1. Definition:

    • Energy released when an isolated gaseous ion is hydrated.

    • Varies with charge density and ionic size.

  2. Hydrated Radius:

    • Distance from the ion’s nucleus to the point at which water molecules remain immobile due to attraction.

Applications of Electronegativity

  1. Predicting acidic and basic properties of metal oxides.

  2. Key Points:

    • Most metal oxides are basic except specific amphoteric ones.

Review Questions

  1. Q: What causes the decrease in atomic radius across d-block elements?

    • A: Increase in nuclear charge.

  2. Q: Why does the atomic radius remain constant in the middle of the transition series?

    • A: Due to poor shielding by d electrons.

  3. Q: What leads to the slight increase in atomic radius towards the end of a transition series?

    • A: Electron pairing in d orbitals.

  4. Q: Which statement is NOT true about d-block elements?

    • A: d electrons shield outer electrons effectively.