In-Depth Study Notes on Chemical Bonding and Periodic Trends

Periodic Trends

  • Atomic radii
  • Ionization energy
  • Electron affinity
  • Metallic character
  • Electronegativity (Oct. 1; review below)

Covalent Bonds

  • Review
  • Energetics: Estimating reaction enthalpies from bond energies

Ionic Bonding

  • Review
  • Energetics: Born-Haber cycle

Electron Affinity

  • Definition:
    • Electron affinity is the energy change when a gaseous atom gains an electron to form a gaseous ion:
      \text{Cl(g) + e}^- \rightarrow \text{Cl}^- \text{(g)}
    • Important Note: This process is NOT the opposite of ionization of Cl(g).
    • Electron affinity is represented as a negative number if the addition of the electron is exothermic (In older books it may be positive).
    • General trend: The more negative the electron affinity (E.A.), the more the atom “likes” to gain an electron.

Electron Affinities Trends

  • General Observations:
    • Electron affinities are generally more negative moving left to right across a row in the periodic table.
    • Electron affinities tend to be generally less negative moving down a column, but the decrease is not substantial.

Increasing Metallic Character

  • Atomic groups demonstrating increasing metallic character:
    • Group 1A (Alkali Metals), Group 2A (Alkaline Earth Metals)
    • Metallic character can be described through properties: malleable, ductile, lustrous, and good conductors of heat and electric current.

Electronegativity

  • Description: A rating system for the affinity of atoms for electrons when participating in a covalent bond.
  • Reference: Page 415 of Tro

Covalent Bonding

  • Characteristics of covalent bonding:
    • Atoms share electrons.
    • Nuclei are attracted to the shared electrons between them.
    • Each pair of shared electrons constitutes one chemical bond.

Interaction Energy of Two Hydrogen Atoms

  • Energy graph illustrating interaction energy with features:
    • Energy scales with distance between the hydrogen atoms.
    • Bond energy can be determined from the graph where energy is released when bonds are formed and required to break bonds.

Energy Dynamics in Bond Formation

  • General concept: Energy is required to break bonds and energy is released when bonds are formed.
  • Example graph with potential energy in kJ/mol showing interaction during bond formation and breakage with respective energy values:
    • Example:
    • Energy released when bond forms: negative bond energy (e.g., -432 kJ/mol)
    • Energy absorbed when bond breaks: positive bond energy (e.g., +432 kJ/mol)

Covalent Bond Energies

  • Average covalent bond energies and bond lengths (in pm):
    • Sample values include:
    • H-H: Bond Energy = 432 kJ/mol, Length = 74 pm
    • N-H: Bond Energy = 391 kJ/mol, Length = 101 pm
    • H-F: Bond Energy = 565 kJ/mol, Length = 92 pm
    • C-H: Bond Energy = 413 kJ/mol, Length = 109 pm

Energy and Chemical Bonds

  • Quote on chemical bonds and energy:
    • "Scientists sometimes say that 'energy is stored in chemical bonds or in a chemical compound,' which may make it sound as if breaking the bonds in the compound releases energy. However, breaking a chemical bond always requires energy. When we say that energy is stored in a compound, or that a compound is energy rich, we mean that the compound can undergo a reaction in which weak bonds break and strong bonds form, thereby releasing energy in the overall process. However, it is always the forming of chemical bonds that releases energy.” (Tro, 6e, pg. 431)

Example Reaction: Formation of Water

  • Reaction:
    • 2H2(g) + O2(g) \rightarrow 2H_2O(g)
    • Energy dynamics:
      1) Break all the bonds to make atoms
      2) Form bonds to make products
  • Breakdown of energy:
    • Energy for breaking bonds (kJ/mol):
    • 2 H-H: +2(432)
    • 1 O=O: +498
    • Energy for making bonds (kJ/mol):
    • 4 H-O: 4(-467)
    • Total:
    • Energy released: -1868 kJ/mol
    • Energy required: +1362 kJ/mol

Calculating Enthalpy Changes

  • Example using average bond energies:
    • Reactions:
    • 2H2(g) + O2(g) \rightarrow 4H(g) + 2O(g)
    • 4H(g) + 2O(g) \rightarrow 2H_2O(g)
  • Total enthalpy change:
    • \Delta H = -506 kJ/mol
    • Literature value: -483.6 kJ/mol

Polarity of Covalent Bonds

  • Types of bonds based on polarity:
    • Nonpolar covalent bonds: Bonds in Cl2, N2, H2 with equal sharing of electrons.
    • Polar covalent bonds: Examples include CO, NH3, H2O, SF6 with unequal sharing of electrons.
    • Ionic bonds: Such as in NaCl, MgCl2, BaO where electrons are transferred, not shared.

Electronegativity and Bonding

  • Electronegativity difference affecting bond types:
    • Zero: Nonpolar covalent
    • Intermediate: Polar covalent
    • Large (2.0 or greater): Ionic
  • Covalent character diminishes as the bond transitions to ionic, while ionic character increases based on electronegativity differences.

Metallic Bonding

  • Definition: Positive nuclei surrounded by a sea of delocalized electrons.
  • Conceptual Differences:
    • Similar to covalent bonding, but with broad delocalized sharing of electrons.