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Lecture 6: Electron Configuration and Chemical Bonding

Electron Configuration

  • Electron Configuration: Specifies how electrons are distributed within orbitals.

    • Representations:
      • Complete specification of quantum numbers.
      • Shorthand specification of orbitals (e.g., 1s2).
      • Orbital diagrams illustrating energy levels and occupancy.

  • Ground State: The most stable electron arrangement.

    • Example: Hydrogen (H) atom (Z = 1)
      • Quantum numbers: n = 1, l = 0, ml = 0, ms = +\frac{1}{2}
      • Electron configuration: 1s1

  • Aufbau Principle: Fill orbitals with electrons starting from the lowest energy level and progressively moving to higher energy levels.

    • Helium (He) atom (Z = 2): 1s2
      • n = 1, l = 0, ml = 0, ms = +\frac{1}{2} and n = 1, l = 0, ml = 0, ms = -\frac{1}{2}
    • Lithium (Li) atom (Z = 3): 1s22s1
      • n = 1, l = 0, ml = 0, ms =+\frac{1}{2}, n = 1, l = 0, ml = 0, ms =-\frac{1}{2} and n = 2, l = 0, ml = 0, ms =+\frac{1}{2}

Filling Rules

  1. Aufbau Principle (Restated): Electrons occupy the most stable available orbital.
  2. Pauli Exclusion Principle: No two electrons can have the same set of four quantum numbers.
  3. Maximum Orbital Capacities:
    • s subshell: 2 electrons
    • p subshell: 6 electrons
    • d subshell: 10 electrons
  4. Energy and n: Higher n value means higher energy.
  5. Energy and l: For equal n, higher l value means higher energy.
  6. Orbital energy is defined by both n and l.

Valence and Core Electrons

  • Valence Electrons: Accessible electrons that determine chemical behavior; located outside filled shells.
  • Core Electrons: Inaccessible electrons located in filled shells.
  • Valence electrons participate in bonding and chemical reactions.
  • Core electrons do not participate in bonding or chemical reactions.
  • Number of valence electrons corresponds to older group numbers (IA, IVA, etc.).
    • Example: Carbon (C) in Group IVA has 4 valence electrons.

Condensed Notation

  • Noble gas configuration represents the element at the end of each period.
  • Saves time by representing core electrons with the symbol of the preceding noble gas.
    • Example: Carbon (C)
      • Shorthand: 1s2 2s2 2p2
      • Condensed notation: [He] 2s2 2p2

Orbital Diagrams

  • Depict orbitals in order of increasing energy (not to scale).
  • Created after writing the shorthand configuration.
    • Example: Lithium (Li): 1s2 2s1

Irregular Configurations and Cations

  • Chromium (Cr) and Copper (Cu) have irregular configurations.
    • Cr: [Ar] 4s1 3d5
    • Cu: [Ar] 4s1 3d10
  • Cations: 4s electrons are removed first.
    • Copper(II) (Cu2+): [Ar] 3d9

Superimposed Orbitals

  • Multi-electron atoms involve the superposition of atomic orbitals populated by electrons.
    • Example: Oxygen (O): 1s2 2s2 2p4

Periodicity

  • Shells generally indicated by the period number (except 3d subshell filled in the 4th period).
  • Groups are characterized by filling distinct subshells (s, p, d).
  • Periodic Table arranges elements in order of increasing atomic number and electron number.
  • Physical and chemical trends are observable.

Atomic Radii

  • Trends:
    • Increases down a group.
    • Decreases across a period.
  • Reasons:
    • Going down a group: Additional electron shell provides more shielding; effective nuclear charge (Z_{eff}) remains approximately constant.
    • Going across a period: No additional shielding; effective nuclear charge (Z_{eff}) increases.

Ionic Radii

  • Cations are smaller than their parent atoms due to electron loss but the same number of protons.
  • Anions are larger than their parent atoms due to additional electrons and the unchanged number of protons.
  • Formation of Group 1 cations and Group 17 anions releases significant energy.
  • Example: Combining Sodium (Na) with Chlorine (Cl2).

Ionization Energy

  • Ionization Energy (I1): Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms.
    • Equation: atom(g) \rightarrow ion^+(g) + e^- I_1 (+ve kJ mol^{-1})
  • Trends:
    • Decreases from top to bottom of the periodic table.
    • Increases from left to right of the periodic table.
  • Trends are inversely related to atomic size.
  • Outermost electron removal requires less energy when farther from the nucleus.
  • I_1 values are large and endothermic (energy must be added).

Higher Ionization Energies

  • Second Ionization Energy (I2): Energy required to remove 1 mole of electrons from 1 mole of gaseous cations.
    • Equation: ion^+(g) \rightarrow ion^{++}(g) + e^- I_2 (+ve kJ mol^{-1})
  • Removing multiple electrons becomes more difficult as cationic charge increases.
  • Noble gas configurations exhibit special stability.

Electron Affinity

  • Electron Affinity (EA1): Energy change associated with adding 1 mole of electrons to 1 mole of gaseous atoms.
    • Equation: atom(g) + e^- \rightarrow ion^-(g) EA_1 (-ve kJ mol^{-1})
  • Trends are less regular.
    • Groups 1A and 2A: low EA_1
    • Groups 6A and 7A: high EA_1
    • Group 8A: low EA_1
  • Low EA_1 means little energy is released (exothermic) or energy may be added (endothermic).

Types of Bonding

  • Metallic bonding: Metal cations fixed in a 'sea' of mobile valence electrons.
  • Ionic bonds: Form between metal cations and non-metal anions.
  • Covalent bonds: Form between two non-metal atoms.

Metallic Bonds

  • Occur between metal atoms.
  • Electron sea model: Delocalized 'sea' of valence electrons around metal ions.

Metal Properties

  • Most exist in a regular but non-rigid metal-ion array with mobile valence electrons.
  • Properties:
    • Solids (liquids).
    • Malleable (can be beaten into sheets), ductile (can be extruded into wires).
    • Good electrical and thermal conductors.
    • Moderate melting points (m.p.) and moderately high boiling points (b.p.).

Ionic Bonds

  • Electrostatic force of attraction between cations and anions in an ionic lattice.
  • Form between low I1 elements and high EA1 elements.

Ionic Solids

  • Generally no discrete molecules.
  • 3D lattice of ions in a fixed arrangement held together by attractive forces between oppositely charged ions.

Ionic Properties

  • Very stable 3D ionic lattice.
  • Properties:
    • Solids: hard, rigid, brittle, high m.p.
    • Low conductivity in solid form.
    • Good conductivity when melted or dissolved.

Covalent Bonds

  • Involve non-metals sharing electron pairs to achieve greater stability.
  • Bond Length: Distance at which molecule has maximum energetic advantage over separated atoms.
  • Bond Energy: Energy required to break the bond (kJ mol-1).

Molecular Covalent Properties

  • Strong covalent intramolecular bonding (inside the molecule).
  • Weak intermolecular bonding (between molecules).
  • Properties:
    • Gases, liquids, or solids.
    • Poor electrical conductors.
    • Generally low melting and boiling points.
    • Solids are soft or brittle.

Network Covalent Solids

  • Strong covalent bonding throughout the structure.
  • No molecules in the structure; 3D array of atoms covalently bound in a lattice.
  • Examples: diamond, graphite, silica (SiO2).
  • Properties:
    • Solids: hard, insoluble.
    • Varying conductivity (insulator or conductor).
    • High melting and boiling points.