Covalent Bonding and Lewis Structures

Introduction to Covalent Bonding

Covalent bonding involves the sharing of electron pairs between atoms, resulting in the formation of molecules. The arrangement of these shared and unshared electrons can be represented using Lewis structures, which are essential for understanding molecular structure and reactivity.

• Covalent Bond: A chemical bond formed by the sharing of one or more pairs of electrons between atoms.

• Lone Pair: A pair of valence electrons not involved in bonding.

• Octet Rule: Atoms tend to form bonds until they are surrounded by eight valence electrons (except hydrogen, which follows the duet rule).

Lewis Structures and the Periodic Table

The number of bonds and lone pairs an atom forms is determined by its group number in the periodic table. Lewis structures help predict molecular geometry and reactivity.

• Group IVA (14): 4 valence electrons, forms 4 bonds (e.g., carbon in CH4).

• Group VA (15): 5 valence electrons, forms 3 bonds (e.g., nitrogen in NH3).

• Group VIA (16): 6 valence electrons, forms 2 bonds (e.g., oxygen in H2O).

• Group VIIA (17): 7 valence electrons, forms 1 bond (e.g., fluorine in HF).

• Group VIIIA (18): 8 valence electrons, forms 0 bonds (noble gases).

Examples of Lewis Structures

• Water (H2O): Oxygen forms two bonds and has two lone pairs.

• Carbon Dioxide (CO2): Carbon forms two double bonds with oxygen, each oxygen has two lone pairs.

• Polyatomic Ions: The total number of electrons is adjusted for the charge (add electrons for negative charge, subtract for positive).

Rules for Lewis Structures:

• Each atom (except H) should have a full octet.

• The total number of electrons must match the sum of valence electrons (adjusted for charge).

• Atoms should have the expected number of bonds; deviations indicate formal charges.

Exceptions to the Octet Rule

Some elements, especially those in period 3 or higher (such as phosphorus and sulfur), can have expanded octets, forming more than four bonds in certain compounds (e.g., PO43-, SO42-).

Three-Dimensional Shapes of Molecules (VSEPR Theory)

VSEPR Theory and Molecular Geometry

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the three-dimensional shapes of molecules based on the repulsion between electron pairs around a central atom. Electron groups (bonded atoms or lone pairs) arrange themselves to minimize repulsion.

• Linear: 2 electron groups, 180° bond angle (e.g., CO2).

• Trigonal Planar: 3 electron groups, 120° bond angle (e.g., SO3).

• Tetrahedral: 4 electron groups, 109.5° bond angle (e.g., CH4).

Polarity of Bonds and Molecules

Bond Polarity and Electronegativity

Bond polarity arises from differences in electronegativity between bonded atoms. The greater the difference, the more polar the bond.

• Nonpolar Covalent: Electronegativity difference < 0.5

• Polar Covalent: Electronegativity difference 0.5–1.9

• Ionic: Electronegativity difference > 1.9

A dipole is created when electrons are unevenly shared, resulting in partial positive and negative charges on the atoms.

Molecular Polarity

The overall polarity of a molecule depends on both the polarity of its bonds and its three-dimensional shape. Symmetrical molecules (e.g., CH4) are nonpolar, while asymmetrical molecules (e.g., H2O, NH3) are polar.

Intermolecular Forces

Types of Intermolecular Forces (IMFs)

Intermolecular forces are attractions between molecules that determine physical properties such as boiling and melting points.

• London Dispersion Forces: Weakest, present in all molecules, especially nonpolar ones.

• Dipole-Dipole Interactions: Occur between polar molecules.

• Hydrogen Bonding: Strongest type of dipole-dipole interaction, occurs when H is bonded to N, O, or F.

States of Matter and Energy Changes

• Endothermic: Absorbs heat (e.g., melting, vaporization).

• Exothermic: Releases heat (e.g., freezing, condensation).

• Enthalpy (ΔH): The heat content change during a process.

Hydrogen Bonding

Definition and Requirements

Hydrogen bonding is a strong intermolecular force that occurs when hydrogen is covalently bonded to a highly electronegative atom (N, O, or F) and is attracted to a lone pair on another electronegative atom.

• Hydrogen Bond Donor: The molecule with the H atom bonded to N, O, or F.

• Hydrogen Bond Acceptor: The molecule with a lone pair on N, O, or F.

Solutions and Their Properties

Definitions and Types of Mixtures

• Mixture: Variable composition, can be homogeneous (solution) or heterogeneous.

• Solution: Homogeneous mixture of ions or small molecules.

• Colloid: Homogeneous mixture with larger particles than a solution.

• Solute: Substance dissolved in a solvent.

• Solubility: Maximum amount of solute that dissolves at a given temperature and pressure.

Solubility and Intermolecular Forces

"Like dissolves like": Polar solvents dissolve polar solutes; nonpolar solvents dissolve nonpolar solutes. Water, being polar, dissolves ionic and polar substances well.

Electrolytes and Osmolarity

• Electrolytes: Substances that conduct electricity in solution by dissociating into ions.

• Strong Electrolytes: Completely dissociate (e.g., NaCl).

• Weak Electrolytes: Partially dissociate (e.g., acetic acid).

• Nonelectrolytes: Do not dissociate (e.g., sucrose).

• Osmolarity: Concentration of all dissolved particles in solution.

Osmosis, Dialysis, and Tonicity

• Osmosis: Movement of solvent through a semipermeable membrane from low to high solute concentration.

• Dialysis: Movement of solvent and small solute particles through a membrane.

• Isotonic Solution: Same osmolarity as cells; no net movement of water.

• Hypotonic Solution: Lower osmolarity than cells; water enters cells, causing swelling.

• Hypertonic Solution: Higher osmolarity than cells; water leaves cells, causing shrinkage.

Summary Table: Bonding and Polarity Patterns