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.