Understanding Covalent Bonds

• Chemical bonds are attractive forces holding atoms together in molecules and compounds.
• These bonds result from interactions between positively charged nuclei and negatively charged electrons.
• Chemical bonding leads to a lower energy state for the atoms, enhancing molecular stability.

Types of Chemical Bonds: Covalent Bonding

• Covalent bonds involve the sharing of electrons to achieve a stable electron configuration.
• This bonding typically occurs between two nonmetal atoms.
• Shared electrons are attracted to both nuclei, effectively bonding the atoms.
• The number of covalent bonds an atom can form is determined by its need to complete its valence shell (octet rule).
• Single bonds share one electron pair, double bonds share two pairs, and triple bonds share three pairs.
• Covalent compounds generally have lower melting and boiling points than ionic compounds due to weaker intermolecular forces.
• Many covalent substances are gases or liquids at room temperature.
• Covalent compounds are poor conductors of electricity due to the absence of freely moving ions or electrons.

Types of Chemical Bonds: Ionic Bonding

• Ionic bonds form through electron transfer, creating ions.
• This bonding occurs typically between a metal and a nonmetal.
• The atom losing electrons becomes a positively charged cation, while the atom gaining electrons becomes a negatively charged anion.
• Electrostatic attraction between oppositely charged ions holds them together in an ionic bond.
• Ionic compounds typically form crystal lattices with repeating three-dimensional ion arrangements.
• Ionic compounds generally have high melting and boiling points because strong electrostatic forces require significant energy to overcome.
• Ionic compounds are good conductors of electricity when dissolved in water or melted due to mobile ions.

Bonding Theories: Valence Bond Theory

• Valence bond theory describes a covalent bond as the overlap of two half-filled atomic orbitals, each containing one electron.
• Overlapping orbitals must be in phase to form a bonding molecular orbital.
• Greater overlap results in a stronger bond.
• Hybridization is the mixing of atomic orbitals to create new hybrid orbitals with different shapes and energies
• Hybridization allows atoms to form stronger and more directional bonds.
• Common types of hybridization include sp, sp2, and sp3 hybridization, each with different geometries.

Bonding Theories: Molecular Orbital Theory

• Molecular orbital theory describes bonding through the combination of atomic orbitals into molecular orbitals that extend over the entire molecule.
• Atomic orbitals combine to form bonding (lower energy) and antibonding (higher energy) molecular orbitals.
• Bonding molecular orbitals are lower in energy than original atomic orbitals, contributing to molecular stability.
• Antibonding molecular orbitals are higher in energy, reducing stability if occupied.
• Bond order is calculated as (number of bonding electrons - number of antibonding electrons) / 2.
• A bond order greater than 0 indicates a stable molecule.
• Molecular orbital diagrams illustrate the relative energies and occupancies of molecular orbitals.

Electronegativity

• Electronegativity measures an atom's ability to attract electrons in a chemical bond.
• Electronegativity values predict bond polarity.
• Fluorine is the most electronegative element.
• Electronegativity generally increases across a period (left to right) and decreases down a group in the periodic table.

Polar vs Nonpolar Bonds

• A polar covalent bond involves unequal electron sharing due to electronegativity differences.
• The more electronegative atom gains a partial negative charge (δ-), and the other atom gains a partial positive charge (δ+).
• A dipole moment measures bond polarity, calculated as the magnitude of partial charges multiplied by the distance between atoms.
• Nonpolar covalent bonds involve equal electron sharing.
• This occurs when bonded atoms have similar electronegativity values.
• Bonds between two atoms of the same element are always nonpolar.
• Electronegativity differences: less than 0.4 indicates a nonpolar covalent bond, between 0.4 and 1.7 indicates a polar covalent bond, and greater than 1.7 indicates an ionic bond.
• Molecular polarity depends on both individual bond polarity and molecular shape.
• Molecules with polar bonds can be nonpolar if bond dipoles cancel out due to symmetry.
• Symmetrical molecules (e.g., CO2, CCl4) can have polar bonds but be nonpolar overall due to canceling dipoles.
• Asymmetrical molecules (e.g., H2O, NH3) have polar bonds and are polar overall because bond dipoles do not cancel.