General Chemistry I: Practice Exam 1 Study Guide

General Chemistry I: Practice Exam 1 Study Guide

Atomic Structure and Isotopes

Atoms are composed of protons, neutrons, and electrons. Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

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

• Average Atomic Mass: The weighted average of the masses of all naturally occurring isotopes of an element.

• Calculation: Multiply the mass of each isotope by its natural abundance (as a decimal), then sum the results.

• Example: If an element has isotopes with masses 10 amu (90%) and 11 amu (10%), the average atomic mass is (10 × 0.90) + (11 × 0.10) = 9 + 1.1 = 10.1 amu.

Chemical Formulas and Empirical Formulas

Chemical formulas represent the types and numbers of atoms in a compound. The empirical formula is the simplest whole-number ratio of atoms in a compound.

• Empirical Formula: The simplest ratio of elements in a compound.

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

• Example: The empirical formula of hydrogen peroxide is CHO, while its molecular formula is C2H2O2.

• Calculation: To determine the empirical formula, divide the percent composition by the atomic masses, then simplify to the smallest whole numbers.

Balancing Chemical Equations

Balancing chemical equations ensures the law of conservation of mass is obeyed. The number of atoms of each element must be the same on both sides of the equation.

• Steps:

  1. Write the unbalanced equation.

  2. Count the number of atoms of each element on both sides.

  3. Add coefficients to balance the atoms.

  4. Check your work.

• Example: $\text{AlCl}_3 (aq) + \text{Ba}_3(\text{PO}_4)_2 (aq) \rightarrow \text{AlPO}_4 (s) + \text{BaCl}_2 (aq)$

Stoichiometry and Limiting Reactants

Stoichiometry involves calculations based on balanced chemical equations. The limiting reactant is the reactant that is completely consumed first, limiting the amount of product formed.

• Steps:

  1. Balance the chemical equation.

  2. Convert given quantities to moles.

  3. Use mole ratios to determine the limiting reactant and amount of product.

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

• Example: If 5.0 g of F2 produces 5.6 g of ClF3, calculate the percent yield using the theoretical yield from stoichiometry.

Molarity and Solution Calculations

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

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

• Titration: Used to determine the concentration of an unknown solution by reacting it with a solution of known concentration.

• Example: If 19.90 mL of 0.250 M NaOH is required to neutralize 30.00 mL of H2SO4, use the balanced equation and molarity formula to find the unknown concentration.

Gas Laws

Gas laws describe the relationships between pressure, volume, temperature, and amount of gas.

• 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 T and P)

• Ideal Gas Law: $PV = nRT$

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

• Example: Calculate the volume a gas will occupy at different conditions using the combined gas law.

Atomic and Molecular Structure

Elements can exist as atoms or molecules. Some elements are diatomic in their natural state.

• Diatomic Elements: H2, N2, O2, F2, Cl2, Br2, I2

• Example: Oxygen (O2) is diatomic at room temperature, while argon (Ar) is monatomic.

Chemical and Physical Changes

Chemical changes result in the formation of new substances, while physical changes do not alter the chemical identity.

• Chemical Change: Rusting of iron, burning of wood, recharging a battery.

• Physical Change: Melting of ice, dissolving sugar in water.

• Example: A color change or gas evolution often indicates a chemical change.

Conversions and Significant Figures

Unit conversions are essential in chemistry. Always use the correct number of significant figures based on the data provided.

• Example: To convert 39.34 yd2 to cm2, use the conversion factors: 1 yd = 0.9144 m, 1 m = 100 cm.

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

Sample Table: Gas Laws Comparison

Law	Equation	Variables Held Constant
Boyle's Law	$P_1V_1 = P_2V_2$	Temperature, amount
Charles's Law	$\frac{V_1}{T_1} = \frac{V_2}{T_2}$	Pressure, amount
Avogadro's Law	$\frac{V_1}{n_1} = \frac{V_2}{n_2}$	Temperature, pressure
Ideal Gas Law	$PV = nRT$	None

Sample Table: Diatomic Elements at Room Temperature

Element	Formula
Hydrogen	H2
Nitrogen	N2
Oxygen	O2
Fluorine	F2
Chlorine	Cl2
Bromine	Br2
Iodine	I2

Additional info:

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

• Tables have been recreated to summarize key comparisons and classifications relevant to the exam content.