CHEM105 Exam 3 Study Guide: Stoichiometry, Solutions, Acids & Bases

Review of Key Concepts from Previous Chapters

SI Units, Prefixes, and Significant Figures

Understanding measurement and calculation accuracy is fundamental in chemistry. SI units and their prefixes allow for conversions across scales, and significant figures ensure precision in reporting results.

• SI Prefixes: Common prefixes include mega (106), kilo (103), centi (10-2), milli (10-3), micro (10-6), nano (10-9).

• Volume Conversion: $1\ \text{mL} = 1\ \text{cm}^3$

• Significant Figures: The number of digits in a measurement that are known with certainty plus one estimated digit. Used to maintain accuracy in calculations.

Diatomic Elements

Seven elements exist naturally as diatomic molecules:

• H2, N2, O2, F2, Cl2, Br2, I2

Ionic vs. Molecular Compounds

• Ionic Compounds: Formed from metals and nonmetals; consist of ions held together by electrostatic forces.

• Molecular Compounds: Formed from nonmetals; consist of molecules held together by covalent bonds.

Empirical Formulas and Naming Compounds

• Empirical Formula: The simplest whole-number ratio of atoms in a compound.

• Naming Ionic Compounds: Name cation first, then anion. Include Roman numerals for transition metals. Polyatomic ions have specific names (e.g., sulfate, nitrate).

• Naming Binary Molecular Compounds: Use prefixes (mono-, di-, tri-, etc.) to indicate the number of each atom.

Chemical Equations and Molar Mass

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

• Molar Mass: The mass of one mole of a substance, calculated from the periodic table.

• Conversions: $\text{moles} = \frac{\text{grams}}{\text{molar mass}}$

Stoichiometry

Relating Reactants and Products

Stoichiometry involves quantitative relationships between reactants and products in a chemical reaction.

• Balanced Equation: Required for stoichiometric calculations.

• Mass-Mole Conversions: Use molar mass and mole ratios from the equation.

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

• Excess Reagent: The reactant left over after the reaction.

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

Solutions and Molarity

Calculating Molarity and Solution Preparation

Molarity is a measure of concentration, defined as moles of solute per liter of solution.

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

• Determining Ion Concentrations: Multiply molarity by the number of ions produced per formula unit.

• Preparing Solutions: Dissolve a known amount of solute in solvent and dilute to a specific volume.

• Using Molarity in Calculations: $\text{moles} = M \times \text{volume (L)}$

• Dilution: $M_1V_1 = M_2V_2$

Example: Preparing a Dilution

• To prepare 100 mL of 0.5 M NaCl from a 2.0 M stock solution: $V_1 = \frac{M_2V_2}{M_1} = \frac{0.5 \times 0.1}{2.0} = 0.025\ \text{L} = 25\ \text{mL}$

Stoichiometry in Solution Reactions

Precipitation and Neutralization Reactions

Stoichiometric calculations can be applied to reactions in solution, such as precipitation and neutralization.

• Precipitation Reaction: Two soluble salts react to form an insoluble product (precipitate).

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

• Use Molarity: Calculate moles of reactants from volume and concentration.

Electrolytes and Dissolution Processes

Classification of Electrolytes

Electrolytes are substances that produce ions in solution, classified as strong, weak, or nonelectrolytes.

• Strong Electrolytes: Completely dissociate into ions (e.g., NaCl, strong acids).

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

• Nonelectrolytes: Do not produce ions (e.g., sugar).

Dissolution Processes

• Strong Acids: Complete ionization: $\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-$

• Weak Acids/Bases: Partial ionization: $\text{CH}_3\text{COOH} \rightleftharpoons \text{H}^+ + \text{CH}_3\text{COO}^-$

• Soluble Ionic Compounds: Dissociation: $\text{NaOH} \rightarrow \text{Na}^+ + \text{OH}^-$

• Nonelectrolytes: Simple dissolution, no ions produced.

Writing Dissolution Equations

• Write complete and balanced equations for dissolution of ionic compounds, acids, bases, and nonelectrolytes.

Precipitation Reactions

Solubility Rules and Double Displacement Reactions

Solubility rules help predict whether a compound will dissolve in water. Precipitation reactions are a type of double displacement (metathesis) reaction.

• Solubility Rules: Used to determine if a product is insoluble and will precipitate.

• Double Displacement Reaction: $\text{AB} + \text{CD} \rightarrow \text{AD} + \text{CB}$

• Types of Equations:

  • Molecular Equation: Shows reactants and products as compounds.

  • Complete Ionic Equation: Shows all soluble ionic substances as ions.

  • Net Ionic Equation: Shows only the ions and molecules directly involved in the reaction.

• Spectator Ions: Ions that do not participate in the reaction.

Acids and Bases

Strong Acids and Bases

Acids and bases are classified by their strength, which refers to their degree of ionization in water.

• Seven Strong Acids:

  • Hydrochloric acid (HCl)

  • Hydrobromic acid (HBr)

  • Hydroiodic acid (HI)

  • Nitric acid (HNO3)

  • Perchloric acid (HClO4)

  • Sulfuric acid (H2SO4)

  • Chloric acid (HClO3)

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

• Acid Identification: Acids typically have H at the front of the formula.

Dissolution of Acids

• Strong Acids: Complete dissociation: $\text{HNO}_3 \rightarrow \text{H}^+ + \text{NO}_3^-$

• Weak Acids: Partial dissociation: $\text{HF} \rightleftharpoons \text{H}^+ + \text{F}^-$

Neutralization Reactions

Predicting Products and Writing Equations

Neutralization reactions occur when an acid reacts with a base to produce water and a salt. These are a type of double displacement reaction.

• General Reaction: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

• Equation Types:

  • Molecular Equation: Shows reactants and products as compounds.

  • Complete Ionic Equation: Shows all soluble ionic substances as ions.

  • Net Ionic Equation: Shows only the ions and molecules directly involved in the reaction.

Example: Neutralization Reaction

• Molecular: $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$

• Complete Ionic: $\text{H}^+ + \text{Cl}^- + \text{Na}^+ + \text{OH}^- \rightarrow \text{Na}^+ + \text{Cl}^- + \text{H}_2\text{O}$

• Net Ionic: $\text{H}^+ + \text{OH}^- \rightarrow \text{H}_2\text{O}$

Summary Table: Electrolyte Classification

Type	Example	Dissolution Process
Strong Electrolyte	NaCl, HCl	Complete dissociation
Weak Electrolyte	CH3COOH	Partial ionization
Nonelectrolyte	Glucose	No ionization

Additional info: Students will be provided with a periodic table and solubility rules for reference during the exam.