General Chemistry Exam and Practice Questions: Stoichiometry, Solutions, Chemical Reactions, and Atomic Structure

Stoichiometry and Chemical Calculations

Mass, Moles, and Empirical Formulas

Stoichiometry involves the calculation of reactants and products in chemical reactions using relationships based on the mole concept and balanced equations.

• Mole: The amount of substance containing Avogadro's number ($6.022 \times 10^{23}$) of entities (atoms, molecules, ions).

• Molar Mass: The mass of one mole of a substance, expressed in grams per mole (g/mol).

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

• Percent Yield: The ratio of actual yield to theoretical yield, expressed as a percentage.

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

Example: If 13g of a compound is produced from a reaction with a theoretical yield of 18g, the percent yield is:

• $\frac{13}{18} \times 100\% = 72.2\%$

Limiting Reactant and Excess Reactant

The limiting reactant is the reactant that is completely consumed first, limiting the amount of product formed. The excess reactant is the one left over after the reaction.

• To determine the limiting reactant:

  1. Calculate the moles of each reactant.

  2. Use the balanced equation to find which reactant produces the least amount of product.

• To find the mass of excess reactant remaining:

  1. Calculate how much of the excess reactant is used.

  2. Subtract from the initial amount.

Example: If 40g of C4H10 is combusted and only 17.6g of CO2 is produced, calculate the percent yield and the mass of excess reactant.

Solutions and Concentrations

Molarity and Dilution

Molarity (M) is the concentration of a solution, defined as moles of solute per liter of solution.

• $M = \frac{\text{moles of solute}}{\text{volume of solution in liters}}$

Dilution: When a solution is diluted, the number of moles of solute remains constant, but the volume increases.

• $M_1V_1 = M_2V_2$

Example: If 100.0 mL of a 1.00 M solution is diluted to 1000.0 mL, the final concentration is:

• $M_2 = \frac{1.00 \times 100.0}{1000.0} = 0.10$ M

Precipitation Reactions

When two aqueous solutions are mixed, a precipitate may form if an insoluble product is produced.

• Example: Mixing AgNO3 and NaCl forms AgCl (s) as a precipitate.

Chemical Reactions and Equations

Types of Chemical Reactions

• Combustion: Reaction of a substance with oxygen to produce energy, CO2, and H2O.

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

• Redox (Oxidation-Reduction): Transfer of electrons between species; oxidation is loss of electrons, reduction is gain.

Example: In the reaction $\text{Fe} + \text{Cu}^{2+} \rightarrow \text{Fe}^{2+} + \text{Cu}$, Fe is oxidized and Cu2+ is reduced.

Balancing Chemical Equations

Balanced equations have equal numbers of each atom on both sides.

• Example: $\text{TiCl}_4 + \text{O}_2 \rightarrow \text{TiO}_2 + 2\text{Cl}_2$

Atomic Structure and Isotopes

Atomic Number, Mass Number, and Isotopes

Atoms are composed of protons, neutrons, and electrons. The atomic number is the number of protons; the mass number is the sum of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Average Atomic Mass: Weighted average of the masses of all isotopes.

• $\text{Average Atomic Mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

Example: Magnesium has three isotopes: $^{24}\text{Mg}$, $^{25}\text{Mg}$, $^{26}\text{Mg}$.

Physical Properties and Density

Density Calculations

Density is mass per unit volume.

• $\text{Density} = \frac{\text{mass}}{\text{volume}}$

Example: If a substance has a mass of 9.83 g and a volume of 1 cm3, its density is 9.83 g/cm3.

Chemical Nomenclature and Classification

Types of Compounds

• Ionic Compounds: Composed of cations and anions (e.g., NaCl, KCl).

• Molecular Compounds: Composed of nonmetals (e.g., H2O, CO2).

• Heteronuclear Diatomic Molecules: Two different atoms (e.g., CO).

• Homonuclear Diatomic Molecules: Two identical atoms (e.g., O2).

Example Table: Classification of Molecules

Type	Example
Heteronuclear diatomic	CO
Homonuclear diatomic	O2
Molecular compound	H2O
Ionic compound	NaCl

Physical and Chemical Changes

Distinguishing Physical and Chemical Changes

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

• Chemical Change: Change that alters the chemical composition (e.g., combustion, rusting).

Example: Burning a steak is a chemical change; melting ice is a physical change.

Sample Calculations and Practice Problems

Representative Problems

• Calculate the number of atoms in a given mass of a compound using Avogadro's number.

• Determine the limiting reactant and calculate the mass of excess reactant remaining.

• Find the molarity after dilution or mixing solutions.

• Balance chemical equations and identify types of reactions.

• Classify compounds and molecules based on their composition.

Additional info: Some questions involve multi-step stoichiometry, empirical formula determination, and solution preparation, which are foundational skills in general chemistry.