Chapter 4 Notes on Structure and Properties of Ionic and Covalent Compounds

4.1 Chemical Bonding

  • Chemical Bond: The force of attraction between two atoms in a compound, primarily involving valence electrons.
Lewis Symbols
  • Lewis Symbol (Lewis Structure): Representation of atoms and their bonds using the element symbol and dots for valence electrons.
Principal Types of Chemical Bonds
  • Ionic Bond:
    • Definition: Transfer of one or more electrons from one atom to another leading to attraction due to opposite charges.
    • Example: Formation of sodium chloride (NaCl).
  • Covalent Bond:
    • Definition: Attractive force due to the sharing of electrons between atoms.
    • Example: Formation of hydrogen (H2) where hydrogen atoms share electrons for a stable configuration.
Formation of Sodium Chloride (NaCl)
  • Sodium (Na):
    • Low ionization energy;
    • Loses an electron to form Na+ and achieves Ne configuration.
  • Chlorine (Cl):
    • High electron affinity;
    • Gains an electron to form Cl- and achieves Ar configuration.
Essential Features of Ionic Bonding
  • Atoms with low ionization energy (I.E.) tend to form positive ions.
  • Atoms with high electron affinity (E.A.) tend to form negative ions.
  • Interactions occur through electron transfer, holding ions together by electrostatic forces.
Covalent Bonding
  • Observed between atoms with similar tendencies to gain/lose electrons.
  • Shared electrons create a covalent bond, providing both atoms with noble gas configurations.
Features of Covalent Bonds
  • Occur between nonmetals;
  • Diatomic molecules (e.g H2, N2) exhibit equal sharing of electrons.
Polar Covalent Bonding
  • Polar Covalent Bond:
    • Unequally shared electron pairs.
    • Electronegativity determines bond polarity.
  • Example: Difference in electronegativity between H-F (1.9) vs H-Cl (0.9).

4.2 Naming Compounds and Writing Formulas

  • Nomenclature: System for naming chemical compounds.
Ionic Compounds
  • Commonly formed by metals (cations) and nonmetals (anions).
  • Example: NaCl as sodium chloride.
  • Crystal Lattice: Regular three-dimensional arrangement of ions.
Writing Formulas of Ionic Compounds
  1. Determine the charge of the ions.
  2. Combine ions to achieve a net charge of zero.
  3. Examples:
    • Sodium sulfide: Na2S
    • Copper(II) carbonate: CuCO3
Naming Ionic Compounds Using the Stock System
  1. Name the cation and anion with appropriate suffixes.
  2. If cation has multiple charges, use Roman numerals (e.g., FeCl3 as iron(III) chloride).
Known Ions
  • Monatomic Ions: Single atom ions like K+ (potassium) and O2- (oxide).
  • Polyatomic Ions: Ions consisting of multiple atoms, e.g., NH4+ (ammonium), SO42- (sulfate).
Common Polyatomic Ions
  • Examples:
    • NO3- (nitrate), OH- (hydroxide), CO32- (carbonate).
Covalent Compounds
  • Formed from nonmetals without extensive three-dimensional structure.
  • Naming includes prefixes for atom numbers (e.g., N2O5 as dinitrogen pentoxide).

4.3 Properties of Ionic and Covalent Compounds

  • Physical States: Ionic compounds are solids; covalent can be solids, liquids, or gases.
  • Melting and Boiling Points:
    • Ionic compounds have higher due to strong ionic bonds.
    • Covalent typically lower.
Solutions and Conductivity
  • Ionic compounds dissolve, dissociating into ions, allowing electrical conductivity (electrolytes).
  • Covalent usually do not dissociate and are non-electrolytes.

4.4 Drawing Lewis Structures

  • Guidelines:
    1. Write skeletal structure with least electronegative atom in center.
    2. Count total valence electrons, adjusting for charges.
    3. Connect central atom to surrounding; satisfy octet.
    4. Adjust bonds if necessary (double/triple bonds).
Examples of Lewis Structures
  • H2O, NH3, CO2.
Bonding Types
  • Single, Double, Triple Bonds: Indicate the number of shared electron pairs.
    • Bond energy decreases from triple to single.
  • Resonance Structures: When multiple valid structures can represent the same molecule.
Exception to the Octet Rule
  • Incomplete Octet: Example - BF3
  • Odd Electron Molecules: Example - NO
  • Expanded Octet: Elements from the 3rd period and below can accommodate more than 8 electrons (e.g., SF6).

4.5 Properties Based on Molecular Geometry

  • Intermolecular Forces: Influential in solubility and boiling/melting points.
Solubility and Miscibility
  • "Like dissolves like": Polar substances dissolve in polar solvents (e.g., NH3 in water).
Boiling and Melting Points
  • Greater intermolecular forces lead to higher boiling/melting points.
  • Polar molecules generally exhibit higher values than nonpolar molecules.
Summary of Conceptual Understanding
  • Understand ionic vs covalent bonding, differentiation in properties, naming conventions, Lewis structures, and molecular geometry principles to determine polarity and solubility.