Chemical bonds are the glue that holds atoms together, forming molecules and materials. Ionic and covalent bonds are the two main types, each with unique properties and formation processes.

Understanding these bonds is crucial for grasping how atoms interact and form compounds. This knowledge forms the foundation for predicting molecular structures, properties, and chemical reactions in various fields of chemistry.

Types of Chemical Bonds

Ionic and Covalent Bonding Fundamentals

• Ionic bonds form between metals and nonmetals through electron transfer
  • Involves complete transfer of electrons from one atom to another
  • Results in oppositely charged ions held together by electrostatic attraction
  • Commonly found in salts (sodium chloride)
• Covalent bonds occur between nonmetals through electron sharing
  • Atoms share one or more pairs of electrons
  • Forms molecules or network structures
  • Found in organic compounds (methane)
• Polar covalent bonds arise when electrons are shared unequally
  • Electron density shifts towards more electronegative atom
  • Creates a partial negative charge on one atom and partial positive on the other
  • Occurs in molecules like water (H2O)
• Nonpolar covalent bonds involve equal sharing of electrons
  • Electron density distributed symmetrically between atoms
  • No charge separation within the molecule
  • Seen in diatomic molecules (oxygen gas, O2)

Bond Polarity and Partial Charges

• Bond polarity measures the degree of electron sharing inequality
  • Determined by difference in electronegativity between bonded atoms
  • Ranges from nonpolar (equal sharing) to highly polar (near-ionic)
  • Affects molecular properties (boiling point, solubility)
• Partial charges develop in polar covalent bonds
  • Denoted as δ+ for slightly positive and δ- for slightly negative
  • Magnitude depends on electronegativity difference
  • Influences intermolecular interactions (hydrogen bonding)
• Dipole moments quantify overall molecular polarity
  • Vector sum of individual bond dipoles
  • Affected by both bond polarity and molecular geometry
  • Measured in Debye units

Factors Influencing Bond Formation

Atomic Properties and Bonding Tendencies

• Electronegativity measures an atom's ability to attract electrons in a bond
  • Increases from left to right across the periodic table
  • Decreases from top to bottom within a group
  • Fluorine has the highest electronegativity (3.98 on Pauling scale)
• Electron affinity quantifies energy change when an atom gains an electron
  • Generally increases from left to right across the periodic table
  • Halogens have high electron affinities
  • Affects an atom's tendency to form anions
• Ionization energy represents the energy required to remove an electron
  • First ionization energy increases across a period
  • Decreases down a group due to increasing atomic size
  • Multiple ionization energies exist for removing subsequent electrons

Energetics of Ionic Compound Formation

• Lattice energy measures the strength of ionic bonds in a crystal
  • Defined as energy released when gaseous ions form a solid
  • Increases with increasing charge and decreasing ion size
  • Contributes to the stability of ionic compounds
• Octet rule guides bonding behavior for main group elements
  • Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons
  • Explains common ionic charges (Na+ Cl-)
  • Has exceptions (expanded octets, incomplete octets)
• Born-Haber cycle relates various energy terms in ionic compound formation
  • Includes ionization energy, electron affinity, and lattice energy
  • Used to calculate enthalpy of formation for ionic compounds
  • Demonstrates the role of energetics in bond formation

Molecular Structure and Geometry

Lewis Structures and Electron Arrangement

• Lewis structures depict valence electrons in molecules and ions
  • Use dots to represent valence electrons
  • Show bonding and nonbonding electron pairs
  • Follow octet rule for main group elements
• Steps to draw Lewis structures
  1. Count total valence electrons
  2. Connect atoms with single bonds
  3. Place remaining electrons as lone pairs
  4. Check octets and adjust if necessary
• Resonance structures represent electron delocalization
  • Multiple valid Lewis structures for a single molecule
  • Actual structure is a hybrid of all resonance forms
  • Indicated by double-headed arrow (nitrate ion, NO3-)

Predicting Molecular Shapes

• Electron-domain geometry based on number of electron domains
  • Includes both bonding and nonbonding electron pairs
  • Determines overall electron distribution around central atom
  • Follows VSEPR (Valence Shell Electron Pair Repulsion) theory
• Molecular geometry describes actual shape of molecule
  • Considers only the arrangement of atoms, not lone pairs
  • Can differ from electron-domain geometry due to lone pairs
  • Affects molecular properties (polarity, reactivity)
• Common molecular geometries
  • Linear (2 domains, CO2)
  • Trigonal planar (3 domains, BF3)
  • Tetrahedral (4 domains, CH4)
  • Trigonal bipyramidal (5 domains, PCl5)
  • Octahedral (6 domains, SF6)
• Bond angles influenced by electron domain repulsions
  • Lone pairs exert stronger repulsion than bonding pairs
  • Results in deviations from ideal geometry (bent shape of H2O)