Chem 140: Final Exam Prep – Comprehensive Study Notes

Chapter 1: Essential Ideas of Chemistry and Science

1.1 The Fundamental Steps in Developing a Model

Chemistry relies on the scientific method to develop models that explain natural phenomena.

• Observation: Gathering data through senses or instruments.

• Hypothesis: Proposing a tentative explanation.

• Experimentation: Testing hypotheses under controlled conditions.

• Theory: A well-substantiated explanation based on repeated experiments.

• Law: A concise statement describing a consistent relationship in nature.

1.2 Units of Measurement

• SI Units: The International System of Units is used for scientific measurements (meter, kilogram, second, mole, kelvin, ampere, candela).

• Prefixes: Used to indicate multiples or fractions of units (e.g., kilo-, milli-).

1.3 Scientific Notation and Significant Figures

• Scientific Notation: Expresses numbers as a product of a coefficient and a power of ten (e.g., $6.02 \times 10^{23}$).

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

1.4 Dimensional Analysis

• Definition: A method for converting between units using conversion factors.

• Example: To convert 5.0 cm to meters: $5.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.050\ \text{m}$

1.5 Accuracy and Precision

• Accuracy: How close a measurement is to the true value.

• Precision: How close repeated measurements are to each other.

Chapter 2: The Components of Matter

2.1 Elements, Compounds, and Mixtures

• Element: A substance that cannot be broken down into simpler substances.

• Compound: A substance composed of two or more elements chemically combined.

• Mixture: A physical blend of two or more substances.

2.2 Atomic Theory and the Structure of the Atom

• Dalton’s Atomic Theory: All matter is composed of atoms; atoms of the same element are identical.

• Subatomic Particles: Protons (+), neutrons (0), electrons (–).

2.3 The Periodic Table

• Organization: Elements are arranged by increasing atomic number.

• Groups/Families: Vertical columns with similar chemical properties.

• Periods: Horizontal rows.

Chapter 3: Stoichiometry of Formulas and Equations

3.1 Chemical Formulas and Nomenclature

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

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

• Naming Compounds: Follows IUPAC rules for ionic and molecular compounds.

3.2 Balancing Chemical Equations

• Law of Conservation of Mass: Matter is neither created nor destroyed in a chemical reaction.

• Balancing: Adjusting coefficients to have equal numbers of each atom on both sides.

3.3 Stoichiometric Calculations

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

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

• Conversions: Grams ↔ Moles ↔ Particles.

Chapter 4: Chemical Reactions

4.1 Types of Chemical Reactions

• Synthesis: $A + B \rightarrow AB$

• Decomposition: $AB \rightarrow A + B$

• Single Displacement: $A + BC \rightarrow AC + B$

• Double Displacement: $AB + CD \rightarrow AD + CB$

• Combustion: Hydrocarbon + $O_2 \rightarrow CO_2 + H_2O$

4.2 Evidence of Chemical Change

• Color change, gas production, precipitate formation, temperature change.

Chapter 5: Gases and Kinetic Molecular Theory

5.1 Properties of Gases

• Compressibility, expandability, low density, diffusion.

5.2 Gas Laws

• Boyle’s Law: $P_1V_1 = P_2V_2$ (at constant T, n)

• Charles’s Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P, n)

• Avogadro’s Law: $\frac{V_1}{n_1} = \frac{V_2}{n_2}$ (at constant P, T)

• Ideal Gas Law: $PV = nRT$

5.3 Kinetic Molecular Theory

• Explains gas behavior based on particle motion and energy.

Chapter 6: Thermochemistry

6.1 Energy Changes in Chemical Reactions

• Exothermic: Releases heat ($\Delta H < 0$).

• Endothermic: Absorbs heat ($\Delta H > 0$).

6.2 First Law of Thermodynamics

• Energy cannot be created or destroyed, only transformed.

• $\Delta E = q + w$ (change in internal energy = heat + work)

Chapter 7: Quantum Theory and Atomic Structure

7.1 Nature of Light

• Light exhibits both wave-like and particle-like properties.

• Wavelength ($\lambda$), frequency ($\nu$), and speed ($c$) are related: $c = \lambda \nu$

7.2 Atomic Spectra and Bohr Model

• Electrons occupy quantized energy levels.

• Emission spectra provide evidence for quantized states.

Chapter 8: Electron Configuration and Chemical Periodicity

8.1 Electron Configuration

• Describes the arrangement of electrons in an atom.

• Aufbau principle, Pauli exclusion principle, Hund’s rule.

8.2 Periodic Trends

• Atomic Radius: Increases down a group, decreases across a period.

• Ionization Energy: Decreases down a group, increases across a period.

• Electronegativity: Tendency to attract electrons; increases across a period.

Chapter 9: Models of Chemical Bonding

9.1 Ionic and Covalent Bonds

• Ionic Bond: Transfer of electrons from metal to nonmetal.

• Covalent Bond: Sharing of electrons between nonmetals.

9.2 Lewis Structures

• Visual representations of valence electrons in molecules.

Chapter 10: The Shapes of Molecules

10.1 VSEPR Theory

• Valence Shell Electron Pair Repulsion (VSEPR) theory predicts molecular shapes based on electron pair repulsion.

• Common shapes: linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Chapter 11: Theories of Covalent Bonding

11.1 Valence Bond Theory and Hybridization

• Atomic orbitals mix to form hybrid orbitals (sp, sp2, sp3).

• Explains molecular geometry and bonding properties.

11.2 Molecular Orbital Theory

• Atomic orbitals combine to form molecular orbitals that are delocalized over the molecule.

Chapter 12: Intermolecular Forces

12.1 Types of Intermolecular Forces

• London Dispersion: Weak, present in all molecules.

• Dipole-Dipole: Between polar molecules.

• Hydrogen Bonding: Strong, between H and N, O, or F.

Chapter 13: Properties of Solutions

13.1 Solution Formation

• Solute: Substance dissolved.

• Solvent: Substance doing the dissolving.

• Concentration Units: Molarity ($M = \frac{\text{mol solute}}{\text{L solution}}$), molality, percent composition.

Chapter 14: Chemical Kinetics

14.1 Reaction Rates

• Rate = change in concentration per unit time.

• Factors: concentration, temperature, catalysts, surface area.

Chapter 15: Chemical Equilibrium

15.1 The Equilibrium Constant

• For $aA + bB \rightleftharpoons cC + dD$, $K = \frac{[C]^c[D]^d}{[A]^a[B]^b}$

• Le Châtelier’s Principle: System shifts to counteract disturbances.

Chapter 16: Acid-Base Equilibria

16.1 Acids and Bases

• Arrhenius: Acids produce $H^+$, bases produce $OH^-$ in water.

• Bronsted-Lowry: Acids donate protons, bases accept protons.

• pH: $pH = -\log[H^+]$

Chapter 17: Thermodynamics

17.1 Spontaneity and Entropy

• Spontaneous Process: Occurs without outside intervention.

• Entropy ($S$): Measure of disorder; increases in spontaneous processes.

• Gibbs Free Energy: $\Delta G = \Delta H - T\Delta S$

Chapter 18: Electrochemistry

18.1 Redox Reactions and Electrochemical Cells

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Galvanic Cell: Converts chemical energy to electrical energy.

• Cell Potential: $E_{cell} = E_{cathode} - E_{anode}$

Chapter 19: Nuclear Chemistry

19.1 Types of Radioactive Decay

• Alpha ($\alpha$) Decay: Emission of a helium nucleus.

• Beta ($\beta$) Decay: Conversion of neutron to proton or vice versa.

• Gamma ($\gamma$) Decay: Emission of high-energy photons.

19.2 Nuclear Reactions

• Fission: Splitting of a heavy nucleus.

• Fusion: Combining of light nuclei.

Additional info: These notes are structured to provide a comprehensive overview of the main topics in a General Chemistry course, following the chapter outline provided in the study materials. For each chapter, key definitions, equations, and examples are included to facilitate exam preparation.