CHEM 111 Exam 3 Study Guide: Concepts and Skills from Chapters 6, 7, 8, and 10

Overview

This study guide summarizes the key concepts and skills covered in Exam 3 for CHEM 111, focusing on Chapters 6 (Chemical Bonding II), 7 (Chemical Reactions and Chemical Quantities), 8 (Introduction to Solutions and Aqueous Reactions), and 10 (Gases). It includes definitions, equations, examples, and tables to support student understanding and exam preparation.

Chapter 6: Chemical Bonding II

Interpreting Energy Interaction Diagrams

• Energy interaction diagram: Shows the energy changes as atoms approach each other to form a bond.

• Bond length: The distance between nuclei at minimum energy.

• Bond energy: The energy required to break a bond between two atoms.

• Example: The energy minimum in a potential energy diagram corresponds to the bond length and bond energy of a molecule.

Valence Bond Theory

• Valence bond theory: Describes bonding as the overlap of atomic orbitals to form covalent bonds.

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals suitable for bonding.

• Types of hybridization: sp, sp2, sp3, sp3d, sp3d2.

• Example: Methane (CH4) has sp3 hybridization, resulting in a tetrahedral geometry.

Sigma and Pi Bonds

• Sigma (σ) bond: Formed by head-on overlap of orbitals; single bonds are sigma bonds.

• Pi (π) bond: Formed by side-on overlap of p orbitals; present in double and triple bonds.

• Example: Ethylene (C2H4) has a double bond consisting of one sigma and one pi bond.

Molecular Orbital Theory

• Molecular orbital theory: Atomic orbitals combine to form molecular orbitals that are delocalized over the molecule.

• Bonding and antibonding orbitals: Bonding orbitals increase electron density between nuclei; antibonding orbitals decrease it.

• Bond order: $(\text{Bond Order}) = \frac{\text{Number of bonding electrons} - \text{Number of antibonding electrons}}{2}$

• Example: O2 has a bond order of 2, indicating a double bond.

Chapter 7: Chemical Reactions and Chemical Quantities

Physical vs. Chemical Changes

• Physical change: Change in state or appearance without altering chemical composition (e.g., melting ice).

• Chemical change: Formation of new substances via chemical reactions (e.g., rusting iron).

• Example: Burning wood is a chemical change; dissolving sugar in water is a physical change.

Balancing Chemical Equations

• Balanced equation: Equal number of atoms of each element on both sides of the equation.

• Example: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

Stoichiometry

• Stoichiometry: Quantitative relationship between reactants and products in a chemical reaction.

• Mole ratio: Ratio of moles of one substance to another in a balanced equation.

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

• Example: If 10 g of product is obtained but 12 g is expected, percent yield is $\frac{10}{12} \times 100\% = 83.3\%$

Chapter 8: Introduction to Solutions and Aqueous Reactions

Solution Concentrations

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

• Example: 0.5 mol NaCl dissolved in 1 L of water yields a 0.5 M solution.

Electrolytes and Nonelectrolytes

• Electrolyte: Substance that conducts electricity when dissolved in water (e.g., NaCl).

• Nonelectrolyte: Substance that does not conduct electricity in solution (e.g., sugar).

• Example: KNO3 is an electrolyte; C6H12O6 is a nonelectrolyte.

Acids and Bases

• Acid: Substance that donates H+ ions in solution.

• Base: Substance that accepts H+ ions or donates OH- ions.

• Example: HCl is an acid; NaOH is a base.

Precipitation, Neutralization, and Redox Reactions

• Precipitation reaction: Formation of an insoluble product (precipitate) from soluble reactants.

• Neutralization reaction: Acid reacts with base to form water and a salt.

• Redox reaction: Transfer of electrons between species; involves changes in oxidation states.

• Oxidation number: Indicates the degree of oxidation of an atom in a compound.

• Example: $\text{Zn} + \text{Cu}^{2+} \rightarrow \text{Zn}^{2+} + \text{Cu}$ (Zn is oxidized, Cu2+ is reduced)

Chapter 10: Gases

Gas Pressure and Units

• Pressure (P): Force exerted per unit area by gas molecules.

• Common units: atm, mmHg, torr, Pa.

• Conversion: $1\text{ atm} = 760\text{ mmHg} = 101.325\text{ kPa}$

Gas Laws

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P and n)

• Avogadro's Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$ (at constant P and T)

• Ideal Gas Law: $PV = nRT$

• Example: Calculate the volume of 1 mol of gas at STP: $V = \frac{nRT}{P}$

Partial Pressure and Dalton's Law

• Dalton's Law: $P_{\text{total}} = P_1 + P_2 + P_3 + \ldots$

• Partial pressure: Pressure exerted by each gas in a mixture.

• Example: In a mixture of O2 and N2, each gas exerts its own partial pressure.

Kinetic Molecular Theory

• Kinetic molecular theory: Explains gas behavior based on motion of particles.

• Root mean square speed: $u_{\text{rms}} = \sqrt{\frac{3RT}{M}}$

• Diffusion: Mixing of gases due to random motion.

• Effusion: Escape of gas through a small hole.

• Graham's Law: $\text{Rate}_1/\text{Rate}_2 = \sqrt{M_2/M_1}$

• Example: Lighter gases effuse faster than heavier gases.

Lab Techniques and Procedures (Labs 7-10)

• Physical vs. chemical change: Ability to distinguish and identify changes in lab experiments.

• Stoichiometric calculations: Calculate coefficients and concentrations from experimental data.

• Precipitation reactions: Use data to identify unknown substances.

• Graphing: Properly label axes and interpret graphs.

Key Tables

Common Polyatomic Ions

Ion	Name	Formula
Ammonium	NH4+	NH4+
Nitrate	NO3-	NO3-
Sulfate	SO42-	SO42-
Phosphate	PO43-	PO43-
Hydroxide	OH-	OH-
Carbonate	CO32-	CO32-

Solubility Rules (Summary)

Compound Type	Solubility
Group 1 and NH4+ salts	Soluble
Nitrates, acetates, perchlorates	Soluble
Chlorides, bromides, iodides	Soluble (except Ag+, Pb2+, Hg22+)
Sulfates	Soluble (except Ba2+, Sr2+, Pb2+)
Carbonates, phosphates, sulfides	Insoluble (except Group 1 and NH4+)

Additional info:

• Students should review "Key Concepts" and "Key Equations and Relationships" from the textbook for each chapter.

• Practice problems and example questions are provided throughout the guide to reinforce understanding.