General Chemistry Exam Study Guide: Chemical Reactions, Stoichiometry, and Solution Chemistry

Ch. 3: Chemical Formulas and Experimental Data

3.10 Determining a Chemical Formula from Experimental Data

This section covers how to deduce empirical and molecular formulas from laboratory data, a foundational skill in general chemistry.

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

• Molecular Formula: The actual number of atoms of each element in a molecule.

• Key Steps:

  1. Obtain mass percentages or masses of each element.

  2. Convert masses to moles using molar mass.

  3. Divide by the smallest number of moles to get the simplest ratio.

  4. Determine molecular formula using molar mass.

• Example: If a compound contains 40% C, 6.7% H, and 53.3% O by mass, convert each to moles and find the empirical formula.

Ch. 4: Chemical Reactions and Stoichiometry

4.1 Climate Change and Combustion of Fossil Fuels

Understanding combustion reactions is important for environmental chemistry, though this section is not emphasized for the exam.

• Combustion Reaction: Hydrocarbon reacts with O2 to produce CO2 and H2O.

4.2 Writing and Balancing Chemical Equations

Balancing chemical equations ensures the law of conservation of mass is obeyed in chemical reactions.

• Balanced Equation: Equal numbers of each atom on both sides of the equation.

• Steps:

  1. Write the unbalanced equation.

  2. Balance atoms one at a time, starting with the most complex molecule.

  3. Check your work.

• Example: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$

4.3 Reaction Stoichiometry

Stoichiometry involves using balanced equations to relate quantities of reactants and products.

• Mole-to-Mole Ratios: Use coefficients from balanced equations to convert between moles of substances.

• Mass-to-Mass Calculations: Convert mass to moles, use mole ratio, then convert back to mass.

• Example: How many grams of water are produced from 16 g of methane?

4.4 Stoichiometric Relationships: Limiting Reactant, Theoretical Yield, % Yield, Excess Reactant

Identifying the limiting reactant and calculating yields are essential for predicting reaction outcomes.

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

• Excess Reactant: The reactant left over after the reaction is complete.

• Example: If 10 g of A reacts with 15 g of B, determine the limiting reactant and percent yield.

4.5 Types and Examples of Chemical Reactions

Chemical reactions are classified by the changes that occur and the substances involved.

• Combustion: Reaction with oxygen, producing CO2 and H2O.

• Alkali Metals: Highly reactive metals in Group 1 of the periodic table.

• Halogens: Reactive nonmetals in Group 17.

• Reactivity Patterns: Alkali metals react vigorously with water; halogens are strong oxidizers.

Ch. 5: Solution Chemistry

5.2 Solution Concentration

Concentration describes the amount of solute dissolved in a given quantity of solvent.

• Solute: Substance dissolved.

• Solvent: Substance doing the dissolving (often water).

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

• Application: Use molarity as a conversion factor in stoichiometry problems.

5.4 Types of Aqueous Solutions and Solubility

Understanding solution types and solubility is crucial for predicting reaction outcomes in water.

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

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

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

• Non-electrolytes: Do not dissociate (e.g., sugar).

• Solubility Terms: Soluble, insoluble, precipitate.

• Predicting Solubility: Use solubility rules to determine if a compound will dissolve.

5.5 Precipitation Reactions

Precipitation reactions occur when two solutions are mixed and an insoluble product forms.

• Precipitate: Solid formed from reaction in solution.

• Solubility Rules: Guidelines for predicting if a compound will be soluble or form a precipitate.

• Balanced Equation: Write and balance the chemical equation for the reaction.

• Example: Mixing AgNO3 and NaCl forms AgCl (precipitate) and NaNO3.

5.6 Representing Aqueous Reactions: Molecular, Complete Ionic, and Net Ionic Equations

Chemists use different types of equations to represent reactions in solution.

• Molecular Equation: Shows all reactants and products as compounds.

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

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

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

• Example: For AgNO3 + NaCl: Net ionic equation is $\text{Ag}^+ (aq) + \text{Cl}^- (aq) \rightarrow \text{AgCl} (s)$

5.7 Acid-Base Reactions

Acid-base reactions are fundamental in chemistry, involving proton transfer between substances.

• Acid: Substance that donates H+ ions.

• Base: Substance that accepts H+ ions.

• Hydronium Ion: $\text{H}_3\text{O}^+$, formed when acids dissolve in water.

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

• Example Equation: $\text{HCl} (aq) + \text{NaOH} (aq) \rightarrow \text{NaCl} (aq) + \text{H}_2\text{O} (l)$

• Titration: Analytical technique to determine concentration of an acid or base.

5.8 Gas Evolution Reactions

Some reactions produce gases as products, which can be identified by bubbling or odor.

• Common Gases: H2S, NH3, CO2, SO2.

• Example: Mixing Na2CO3 and HCl produces CO2 gas.

Summary Table: Types of Chemical Equations

Type	Description	Example
Molecular Equation	Shows all reactants and products as compounds	AgNO3 (aq) + NaCl (aq) → AgCl (s) + NaNO3 (aq)
Complete Ionic Equation	Shows all strong electrolytes as ions	Ag+ (aq) + NO3- (aq) + Na+ (aq) + Cl- (aq) → AgCl (s) + Na+ (aq) + NO3- (aq)
Net Ionic Equation	Shows only the species that change during the reaction	Ag+ (aq) + Cl- (aq) → AgCl (s)

Exam Preparation Tips

• Use a non-programmable, non-graphing calculator for calculations.

• Bring a number 2 pencil and periodic table; no other aids allowed.

• Practice balancing equations, stoichiometry, and identifying reaction types.

• Memorize solubility rules and strong/weak acid-base definitions.

Additional info: Some sections (e.g., climate change) are noted as not emphasized for the exam. The study guide is structured to help students focus on key General Chemistry concepts likely to be tested.