CHM577: Inorganic Chemistry Study Notes

CHM577: Inorganic Chemistry Notes

Course Overview

  • Instructor: Dr. Amalina Mohd Tajuddin

  • Topic: The Chemistry of the Elements

  • Objective: Understanding periodic classification and electron configuration of elements, and identifying periodic variation in physical properties.

Chapter Outline

  • Periodic classification of elements

  • Construction of electron configurations

  • Identification of periodic variations in physical properties:
      - Effective nuclear charge (Z_eff)
      - Atomic and ionic radii
      - Ionization energy (IE)
      - Electron affinity (EA)

Introduction

  • Dmitri Mendeleev is credited with development of periodic classification of elements based on properties, which was a significant achievement in the 19th century.

  • Ground state electron configurations of elements can be represented by the configurations like ns², ns²np1, etc., where n is the principal quantum number of the outermost subshell.

Classification of Elements

  • Main Groups: (Group 1A to 7A)
      - Known as representative/main group elements
      - Have incompletely filled s or p subshells

  • Group 8A: (except Helium)
      - Have completely filled p subshells
      - Example configurations:
        - He: 1s²
        - Noble gases (e.g., Ne: 1s² 2s² 2p⁶)

  • d-Block Transition Elements (Group 3B to 8B):
      - Have incompletely filled d subshells
      - They form cations through these d subshells.

  • f-Block Transition Elements (Group 4F & 5F):
      - Includes Lanthanides and Actinides, with incompletely filled f subshells

Ground State Electronic Configurations

  • The Aufbau Principle:
      - Principle of arrangement of electrons in orbitals:
        - Electrons fill the lowest energy orbitals before higher ones.
        - This is coupled with Hund's rules and the Pauli exclusion principle:
          - Hund’s First Rule: Electrons fill all orbitals singly before pairing. Unpaired electrons have parallel spins.
          - Pauli Exclusion Principle: No two electrons can have the same set of quantum numbers (n, l, m_l, m_s).

Electron Configurations

  • Examples of Ground-State Electron Configurations for Cations and Anions:
      - Sodium (Na): [Ne]3s¹ ⟶ Na⁺: [Ne]
      - Calcium (Ca): [Ar]4s² ⟶ Ca²⁺: [Ar]
      - Aluminum (Al): [Ne]3s² 3p¹ ⟶ Al³⁺: [Ne]
      - Fluorine (F): 1s² 2s² 2p⁵ ⟶ F⁻: 1s² 2s² 2p⁶ or [Ne]
      - Oxygen (O): 1s² 2s² 2p⁴ ⟶ O²⁻: 1s² 2s² 2p⁶ or [Ne]
      - Nitrogen (N): 1s² 2s² 2p³ ⟶ N³⁻: 1s² 2s² 2p⁶ or [Ne]

  • Cations (positively charged) tend to lose electrons to achieve noble-gas configurations. Anions (negatively charged) gain electrons to achieve the same.

  • Isoelectronic Species:
      - Species with the same electron configuration. Examples:
        - Na⁻, Al³⁺, F⁻, O²⁻, N³⁻, all are isoelectronic with Ne.
        - H⁻ is isoelectronic with He.

Electron Configurations of Transition Metal Cations

  • When forming cations, transition metal atoms lose electrons from the ns orbital first, then from the (n-1)d orbitals:
      - Iron (Fe):
        - Neutral: [Ar]4s² 3d⁶
        - Fe²⁺: [Ar]3d⁶ (removes two from 4s)
        - Fe³⁺: [Ar]3d⁵ (removes from 4s and one from 3d)
      - Manganese (Mn):
        - Neutral: [Ar]4s² 3d⁵
        - Mn²⁺: [Ar]3d⁵

Periodic Table Trends

  • Influenced by three factors:
      1. Energy Level: Electrons in higher energy levels are further from the nucleus.
      2. Charge on the Nucleus: More protons result in a stronger attraction.
      3. Shielding Effect: Depends on the arrangement of electrons in inner shells, affecting how outer electrons feel the nuclear charge.

Atomic Size

  • Atomic Radius Definition: Half the distance between two nuclei in a diatomic molecule.

  • Trends:
      - Group Trends: Atomic radius increases down a group due to increased filled electron shells pushing outer electrons further from the nucleus.
        - Example: H < Li < Na < K < Rb   - Period Trends: Atomic radius decreases across a period from left to right as protons and effective nuclear charge increase, pulling electrons closer to the nucleus.     - Example: Na > Mg > Al > Si > P > S > Cl > Ar

Ionic Size

  • Anions: Larger than their corresponding neutral atoms due to increased electron-electron repulsion; less pull by protons because of added electrons.

  • Cations: Smaller than their corresponding neutral atoms due to reduced electron-electron repulsion; more protons pull fewer electrons closer.

Ionization Energy (IE)

  • Definition: The minimum energy required to remove an electron from a gaseous atom in its ground state.
      - Processes:
        1. I1 + X(g) → X⁺(g) + e⁻ (First Ionization)
        2. I2 + X(g) → X²⁺(g) + e⁻ (Second Ionization)
        3. I3 + X(g) → X³⁺(g) + e⁻ (Third Ionization)

  • Trends:
      1. Group Trends: Ionization energy decreases as you move down because of increased distance and shielding.
      2. Period Trends: Ionization energy increases across a period due to decreasing atomic radius and increasing effective nuclear charge.

Factors Affecting Ionization Energy

  • Greater nuclear charge increases IE.

  • Increased distance from nucleus decreases IE.

  • Filled and half-filled orbitals are lower in energy, making them easier to achieve and thus lower their IE.

  • The shielding effect reduces effective nuclear charge felt by outermost electrons.

Electron Affinity (EA)

  • Definition: The energy change when an electron is added to an atom in the gas phase, forming an anion.

  • Notation:
      - For example:
        - X(g) + e⁻ → X⁻(g)

  • Trends:
      1. Group Trends: EA decreases down a group.
      2. Period Trends: EA increases across a period.

Exceptions in Electron Affinity Trends

  • First-period nonmetals have lower affinities than their counterparts below in the same group due to stability of existing electron configurations.

  • Elements with completely filled or half-filled orbitals have lower electron affinities, e.g., Be, N, Ne, as they tend to retain their stability instead of gaining an electron.