Valence Bond Theory and Molecular Orbital Theory

Hybridization and Molecular Shape

The three-dimensional shape of molecules is determined by the overlap of atomic orbitals, a concept known as hybridization. Hybridization explains how atomic orbitals mix to form new, equivalent hybrid orbitals suitable for the pairing of electrons to form chemical bonds in molecules.

• Hybridization: The process of combining atomic orbitals (s, p, d) to form new hybrid orbitals (sp, sp2, sp3, sp3d, sp3d2).

• Electronic Geometries and Hybridization:

  • Linear: sp

  • Trigonal planar: sp2

  • Tetrahedral: sp3

  • Trigonal bipyramidal: sp3d

  • Octahedral: sp3d2

• Relation to Orbital Diagrams: The number of electron domains around the central atom determines the type of hybridization.

Example: Methane (CH4) has a tetrahedral geometry and uses sp3 hybridization.

Limitations of Hybridization and Introduction to Molecular Orbital Theory

While hybridization explains many molecular shapes, it cannot account for all observed properties, such as magnetism in O2. This leads to the Molecular Orbital (MO) Theory, which describes electrons in molecules as occupying molecular orbitals that extend over the entire molecule.

• Bonding vs. Antibonding Orbitals: Bonding orbitals are lower in energy and stabilize the molecule; antibonding orbitals are higher in energy and can destabilize the molecule.

• Bond Order: Indicates the strength and stability of a bond. Calculated as:

• Filling Molecular Orbitals: Electrons fill molecular orbitals in order of increasing energy, following the Pauli exclusion principle and Hund's rule.

• Paramagnetic vs. Diamagnetic:

  • Paramagnetic: Molecules with unpaired electrons (attracted to a magnetic field).

  • Diamagnetic: Molecules with all electrons paired (repelled by a magnetic field).

• Sigma (σ) and Pi (π) Bonds: Sigma bonds are formed by head-on overlap; pi bonds by side-on overlap. Single bonds are always sigma; double bonds have one sigma and one pi; triple bonds have one sigma and two pi bonds.

Example: In O2, the MO diagram predicts two unpaired electrons, explaining its paramagnetism.

Room Temperature Reactions

Chemical Equations and Balancing

Chemical equations represent chemical reactions, showing reactants and products, their phases, and the direction of the reaction.

• Symbols: (s) solid, (l) liquid, (g) gas, (aq) aqueous; arrows indicate the direction of the reaction.

• Reactants: Listed on the left; Products: on the right.

• Balancing Equations: The number of atoms of each element must be equal on both sides.

Example:

Stoichiometry and Limiting Reactants

• Mole Ratios: Derived from coefficients in balanced equations; used to convert between moles, mass, and number of molecules.

• Writing Equations: Translate word descriptions into chemical equations and balance them.

• Limiting Reactant: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: Maximum amount of product possible from given reactants.

• Excess Reactant: The reactant that remains after the reaction is complete.

• Percent Yield: Measures reaction efficiency:

• Stoichiometric Calculations: Use mole ratios to solve for unknown quantities in reactions.

Example: If 10.0 g of A reacts with 15.0 g of B, determine the limiting reactant and percent yield if 8.0 g of product is obtained.

Types of Chemical Reactions

• Combination (Synthesis): Two or more substances form one product.

• Decomposition: One substance breaks down into two or more products.

• Combustion: A substance reacts with O2, producing heat and light, often forming CO2 and H2O.

• Double-Replacement: Exchange of ions between two compounds; includes precipitation, acid/base neutralization, and gas-forming reactions.

• Redox Reactions: Involve transfer of electrons; oxidation and reduction occur simultaneously.

Example: (precipitation reaction)

Aqueous Solutions and Reactions

Solutions: Solute, Solvent, and Concentration

A solution consists of a solute (substance dissolved) and a solvent (substance doing the dissolving, usually present in greater amount).

• Dilute Solution: Contains a small amount of solute relative to solvent.

• Concentrated Solution: Contains a large amount of solute relative to solvent.

• Molarity (M): The most common unit of concentration, defined as:

• Dilution Formula: Used when diluting a solution:

Where and are the initial concentration and volume, and and are the final concentration and volume.

Example: To prepare 250 mL of 0.10 M NaCl from a 1.0 M stock solution, use the dilution formula to find the required volume of stock solution.

Dissolution, Electrolytes, and Conductivity

• Dissolution in Water: Ionic compounds dissociate into ions; molecular compounds may not.

• Electrolytes: Substances that conduct electricity when dissolved in water.

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

• Weak Electrolytes: Partially dissociate (e.g., acetic acid, NH3).

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

Example: NaCl is a strong electrolyte; C6H12O6 (glucose) is a nonelectrolyte.

Precipitation, Acid-Base, and Gas-Forming Reactions

• Precipitation Reactions: Two aqueous solutions form an insoluble solid (precipitate).

• Solubility Rules: Guidelines to predict if a compound is soluble or insoluble in water.

• Double Displacement Reactions: Exchange of ions; may result in precipitation, gas formation, or neutralization.

• Gas-Forming Reactions: Certain combinations produce a gas (e.g., CO2, H2S).

Example:

Ionic Equations and Motivation for Reactions

• Complete Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only the species that actually change during the reaction.

Example: For :

• Complete ionic:

• Net ionic:

Acids, Bases, and Their Reactions

• Arrhenius Definition: Acids produce H+ in water; bases produce OH-.

• Brønsted-Lowry Definition: Acids are proton donors; bases are proton acceptors.

• Acid Formulas: Binary acids start with H (e.g., HCl); oxyacids contain H, O, and another element (e.g., H2SO4); carboxylic acids end with COOH.

• Naming Acids: Binary acids: "hydro-" + root + "-ic acid"; oxyacids: based on polyatomic ion name.

• Strong Acids: HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4

• Strong Bases: Group 1 and 2 metal hydroxides (e.g., NaOH, KOH, Ca(OH)2).

• Weak Base: Ammonia (NH3).

• Dissociation Equations: Show strong acids and bases dissociating completely in water.

Example:

Stoichiometry in Solution

• Use molarity and volume to calculate moles of solute.

• Apply stoichiometry to determine mass, volume, or concentration in reactions involving solutions.

Example: How many grams of NaCl are needed to prepare 500 mL of 0.20 M solution?

Oxidation Numbers and Redox Reactions

• Assigning Oxidation Numbers: Rules are used to assign oxidation states to elements in compounds and ions.

• Oxidation: Loss of electrons (increase in oxidation number).

• Reduction: Gain of electrons (decrease in oxidation number).

• Oxidizing Agent: Causes oxidation; is reduced.

• Reducing Agent: Causes reduction; is oxidized.

Example: In , Na is oxidized (0 to +1), Cl is reduced (0 to -1).

Additional info: This guide is based on a study checklist for Exam 3 covering Chapters 6-8, focusing on bonding theories, chemical reactions, and aqueous solution chemistry. For detailed molecular orbital diagrams and solubility rules, refer to your textbook or class handouts.

Example of Hybridization of CO2

Each carbon-oxygen bond is composed of one σ bond and one π bond. The σ bonds around carbon involve the end-to-end overlap of its two hybridized sp orbitals with the molecular orbitals of the terminal oxygen atoms. The two remaining p orbitals of carbon, which are used to form each π bond, exist on planes that are perpendicular to each other. Therefore, the planes that encompass the two C=O bonds are also perpendicular to each other.