General Chemistry Midterm Review: Key Concepts and Practice Problems

General Chemistry Midterm Review

Introduction to Chemistry

This section covers the foundational concepts of chemistry, including the classification of matter, physical and chemical changes, and the structure of atoms.

• Homogeneous vs. Heterogeneous Mixtures: - Homogeneous mixture: A mixture with a uniform composition throughout (e.g., saltwater). - Heterogeneous mixture: A mixture with visibly different components or phases (e.g., salad, sand in water).

• Physical vs. Chemical Changes: - Physical change: A change that does not alter the chemical identity of a substance (e.g., melting ice). - Chemical change: A change that results in the formation of new substances (e.g., rusting iron).

Measurement and Significant Figures

Accurate measurement and proper use of significant figures are essential in chemistry calculations.

• Significant Figures: The digits in a measurement that are known with certainty plus one estimated digit. - Example: 441.5 kg (4 significant figures), 1.00250 cm (6 significant figures), 0.00309 kg (3 significant figures).

• Unit Conversions: Use conversion factors to change units (e.g., 12.56 in to cm; 1 in = 2.54 cm).

Density and Experimental Error

Density is a fundamental property of matter, and understanding experimental error is crucial for evaluating results.

• Density Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Percent Error: $\text{Percent Error} = \left| \frac{\text{Experimental Value} - \text{Accepted Value}}{\text{Accepted Value}} \right| \times 100\%$

Atomic Structure

Atoms are composed of protons, neutrons, and electrons. Atomic number and mass number are key identifiers.

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Total number of protons and neutrons.

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

• Subatomic Particles: Protons (positive), neutrons (neutral), electrons (negative).

Atomic Theories

Key historical models of the atom include Dalton's, Rutherford's, and Thomson's theories.

• Dalton: Atoms are indivisible particles; each element consists of identical atoms.

• Thomson: Discovered the electron; proposed the "plum pudding" model.

• Rutherford: Discovered the nucleus via gold foil experiment; atoms are mostly empty space.

Average Atomic Mass

The average atomic mass is calculated using the masses and relative abundances of isotopes.

• Formula: $\text{Average Atomic Mass} = \sum (\text{Isotope Mass} \times \text{Fractional Abundance})$

• Example: For carbon with isotopes $^{12}$C (98.89%) and $^{13}$C (1.11%), calculate the weighted average.

Chemical Symbols and Formulas

Chemical symbols represent elements; formulas represent compounds. Ionic charges and compound formation are based on periodic trends.

• Isotope Symbol: $^{A}_{Z}\text{X}$ (e.g., $^{37}_{17}\text{Cl}$ for chlorine-37).

• Writing Formulas: Combine ions to form neutral compounds (e.g., potassium dichromate: K2Cr2O7).

• Common Polyatomic Ions: Nitrate (NO3-), sulfate (SO42-), phosphate (PO43-).

Chemical Reactions and Equations

Chemical reactions are classified by the type of change and must be balanced to obey the law of conservation of mass.

• Types of Reactions:

  • Synthesis (Combination)

  • Decomposition

  • Single Replacement

  • Double Replacement

  • Combustion

• Balancing Equations: Adjust coefficients to ensure equal numbers of each atom on both sides.

• Example: $\text{HCl} + \text{Ca(OH)}_2 \rightarrow$ products (predict and balance).

Stoichiometry

Stoichiometry involves calculations based on balanced chemical equations, including limiting reactants and percent yield.

• Mole Concept: 1 mole = $6.022 \times 10^{23}$ particles.

• Molar Mass: The mass of one mole of a substance (g/mol).

• Percent Composition: $\text{Percent by Mass} = \frac{\text{Mass of Element}}{\text{Total Mass}} \times 100\%$

• Empirical and Molecular Formulas: Empirical is the simplest ratio; molecular is the actual number of atoms.

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

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

Gases and Gas Laws

Gas behavior is described by several laws relating pressure, volume, temperature, and amount.

• Ideal Gas Law: $PV = nRT$ Where P = pressure, V = volume, n = moles, R = gas constant, T = temperature (K).

• 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)

• Gas Stoichiometry: Use molar volume at STP (22.4 L/mol) for conversions.

Sample Table: Types of Chemical Reactions

Equation	Type of Reaction
H2(g) + Cl2(g) → 2HCl(g)	Synthesis (Combination)
HgO(s) → Hg(l) + O2(g)	Decomposition
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)	Single Replacement
Na2SO4(aq) + Ba(NO3)2(aq) → 2NaNO3(aq) + BaSO4(s)	Double Replacement
C2H6(g) + 7/2 O2(g) → 2CO2(g) + 3H2O(g)	Combustion

Applications and Problem Solving

• Apply concepts to solve problems involving density, percent composition, empirical formulas, and gas laws.

• Interpret and balance chemical equations, predict products, and classify reaction types.

• Calculate limiting reactants, theoretical and percent yields, and perform stoichiometric conversions.

Additional info: These notes are based on a midterm review question set and cover core topics from introductory general chemistry, including measurement, atomic structure, chemical reactions, stoichiometry, and gas laws. For each topic, students should practice applying formulas and solving representative problems as shown in the review questions.