Chemical Bonding Concepts

Chemical Bonding

• Chemical bonds are classified into three types based on the atoms involved: ionic, covalent, and metallic.
• Ionic bonds are formed when metal atoms bond to nonmetal atoms. Electrons are transferred.
• Covalent bonds are formed when nonmetal atoms bond together. Electrons are shared.
• Metallic bonds are formed between metal and metal atoms. Electrons are pooled.
• The atoms are held together by intermolecular forces.

Ionic Bond

• Ionic bonds form when electrons are transferred between atoms.
• This transfer creates oppositely charged ions (anions and cations) that attract each other.
• Ionic bonds are typically formed between metals (which lose electrons) and nonmetals (which gain electrons).
• The resulting ions arrange themselves in a crystal lattice structure.
• Examples of ionic compounds include NaCl (sodium chloride).

Covalent Bond

• Covalent bonds form when atoms share electrons.
• Each atom contributes one or more electrons to the shared pair(s).
• Covalent bonds are typically formed between nonmetals.
• Shared electrons hold the atoms together because they are attracted to the nuclei of both atoms.
• Single covalent bonds involve one electron pair.
• Double covalent bonds involve two electron pairs.
• Triple covalent bonds involve three electron pairs.
• Examples of covalent compounds include H₂O (water) and CH₄ (methane).

Coordinate Bond

• A coordinate bond is a type of covalent bond where both electrons in the shared pair come from the same atom.
• The "donor" atom provides both of the electrons and the other atom (acceptor) uses the electron pair.
• Acid is the acceptor electron, and base is the donor electron.
• Example H+ + NH3 = NH₄+

Lewis Structures

• Lewis structures are diagrams that show the valence electrons of atoms in a molecule.
• Valence electrons are represented by dots around the atomic symbol.
• The number of valence electrons for main-group elements is equal to the group number of the element.
• Dots are first arranged singly before pairing (exceptions for helium).

Intermolecular Forces

• Intermolecular forces (IMFs) are attractive forces between molecules.
• These forces are generally weaker than intramolecular forces (forces within a molecule).
• IMF strength affects physical properties like melting/boiling points, solubility, and rates of chemical reactions.
• Types of IMF include:
  • London dispersion forces (weakest): present in all molecules
  • Dipole-dipole forces (moderate): present in polar molecules.
  • Hydrogen bonds (strongest): present in molecules with O-H, N-H, or H-F bonds.

Ion-dipole Forces

• Ion-dipole forces are electrostatic attractions between a charged ion and a polar molecule.
• They are a key factor in the dissolution of ionic compounds in polar solvents such as water.
• The strength of these forces is influenced by both the magnitude of the ion charge and the polarity of the molecule.

Electronegativity

• Electronegativity (EN) is an atom's ability to attract electrons in a chemical bond.
• EN values generally increase across a row of the periodic table and decrease down a column.
• Differences in electronegativity between atoms in a bond determine bond type (polar or nonpolar).

Bond Types Based on Electronegativity

• A small difference (0.0-0.4) in electronegativity represents a nonpolar covalent bond.
• An intermediate difference (0.4-2.0) in electronegativity represents a polar covalent bond.
• A large difference (≥2.0) in electronegativity represents an ionic bond.