General Chemistry Exam 2 Study Guide: Lewis Structures, Bonding, and Chemical Reactions

Exam Overview

This study guide covers the key topics for Exam 2 in a General Chemistry course, focusing on Lewis structures, bonding, molecular geometry, chemical reactions, and solution chemistry. The exam includes multiple-choice, fill-in-the-blank, numerical-answer, and short-answer questions. Students are permitted a reference sheet, scratch paper, and a scientific calculator.

Bonding and Lewis Structures

Types of Chemical Bonds

• Pure Covalent Bonds: Electrons are shared equally between atoms (e.g., H2, Cl2).

• Polar Covalent Bonds: Electrons are shared unequally, resulting in partial charges (e.g., H2O).

• Ionic Bonds: Electrons are transferred from one atom to another, forming ions (e.g., NaCl).

Example: In HCl, the bond is polar covalent because chlorine is more electronegative than hydrogen.

Lewis Structures

• Definition: Diagrams that show the bonding between atoms and the lone pairs of electrons in a molecule.

• Steps to Draw Lewis Structures:

  1. Count total valence electrons.

  2. Arrange atoms (central atom is usually least electronegative).

  3. Connect atoms with single bonds.

  4. Distribute remaining electrons to complete octets (or duets for H).

  5. Form double/triple bonds if necessary to satisfy the octet rule.

• Use of Lewis Symbols: Represent valence electrons as dots around element symbols.

Example: The Lewis structure for CO2 is O=C=O, with each O atom having two lone pairs.

Resonance and Formal Charge

• Resonance Structures: Multiple valid Lewis structures for a molecule, differing only in electron placement.

• Formal Charge: Calculated as: $\text{Formal Charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$

• Choose the structure with the lowest formal charges and negative charges on the most electronegative atoms.

Exceptions to the Octet Rule

• Some molecules have atoms with fewer or more than eight electrons (e.g., BF3, SF6).

Molecular Geometry and Bonding Theories

VSEPR Theory

• Valence Shell Electron Pair Repulsion (VSEPR) Theory: Predicts molecular shapes based on electron pair repulsion.

• Common Geometries:

  • Linear: 180° (e.g., CO2)

  • Trigonal planar: 120° (e.g., BF3)

  • Tetrahedral: 109.5° (e.g., CH4)

  • Trigonal bipyramidal: 90°, 120° (e.g., PCl5)

  • Octahedral: 90° (e.g., SF6)

Bond Polarity and Dipole Moments

• Bond Polarity: Determined by the difference in electronegativity between bonded atoms.

• Dipole Moment: A measure of molecular polarity; calculated as: $\mu = q \times r$ where $q$ is the magnitude of the charge and $r$ is the distance between charges.

Hybridization and Bond Order

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals (e.g., sp, sp2, sp3).

• Bond Order: Number of chemical bonds between a pair of atoms. $\text{Bond Order} = \frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2}$

Chemical Reactions and Stoichiometry

Types of Chemical Reactions

• Physical vs. Chemical Change: Physical changes do not alter chemical composition; chemical changes do.

• Balancing Chemical Equations: Ensure the same number of each atom on both sides of the equation.

• Combustion Reactions: Hydrocarbon reacts with O2 to form CO2 and H2O.

Stoichiometry

• Mole Concept: 1 mole = $6.022 \times 10^{23}$ particles.

• Mole Ratios: Use coefficients from balanced equations to relate amounts of reactants and products.

• 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.

• Percent Yield: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Solutions and Solution Chemistry

Concentration Units

• Molarity (M): $\text{Molarity} = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Solution Dilution: $M_1V_1 = M_2V_2$

Electrolytes and Solubility

• Electrolytes: Substances that conduct electricity when dissolved in water (strong, weak, nonelectrolytes).

• Solubility Guidelines: Rules to predict if an ionic compound is soluble in water.

• Precipitation Reactions: Occur when two solutions mix to form an insoluble product (precipitate).

Acids, Bases, and Neutralization

• Acids: Proton donors; Bases: Proton acceptors.

• Neutralization Reaction: Acid + Base → Salt + Water.

• Net Ionic Equations: Show only the species that change during the reaction.

Titration

• Titration: Technique to determine concentration of a solution using a solution of known concentration.

• Equivalence Point: Point at which moles of acid equal moles of base.

Redox Reactions

Oxidation and Reduction

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Oxidizing Agent: Causes oxidation (is reduced).

• Reducing Agent: Causes reduction (is oxidized).

Example: In the reaction Zn + Cu2+ → Zn2+ + Cu, Zn is oxidized and Cu2+ is reduced.

Summary Table: Key Concepts and Formulas

Concept	Key Formula/Definition
Formal Charge	$\text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$
Molarity	$\text{Molarity} = \frac{\text{moles of solute}}{\text{liters of solution}}$
Percent Yield	$\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$
Bond Order	$\text{Bond Order} = \frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2}$
Solution Dilution	$M_1V_1 = M_2V_2$

Additional info:

• Some context and explanations have been expanded for clarity and completeness.

• Specific examples and formulas have been added to ensure the guide is self-contained and exam-ready.