General Chemistry Final Exam Study Guide

General Chemistry Final Exam Study Guide

Introduction to Chemistry

This section covers foundational concepts in chemistry, including definitions, classifications, and basic properties of matter.

• Matter: Anything that has mass and occupies space.

• Atoms, Elements, Compounds: Atoms are the basic units of matter; elements are pure substances consisting of one type of atom; compounds are substances formed from two or more elements chemically bonded.

• States of Matter: Solid, liquid, and gas; differences include shape, volume, and particle arrangement.

• Mixtures vs. Pure Substances: Mixtures contain two or more substances physically combined; pure substances have a fixed composition.

• Homogeneous vs. Heterogeneous Mixtures: Homogeneous mixtures have uniform composition; heterogeneous mixtures do not.

• Physical vs. Chemical Properties: Physical properties can be observed without changing the substance; chemical properties describe a substance's ability to undergo chemical changes.

Measurement and Units

Understanding measurement systems and conversions is essential for quantitative chemistry.

• SI Units: Standard units for length (meter), mass (kilogram), time (second), temperature (kelvin), amount of substance (mole).

• Unit Conversion: Use conversion factors to change between units.

• Temperature Scales: Celsius, Kelvin, Fahrenheit; conversion formulas:

  • $K = ^\circ C + 273.15$

  • $^\circ F = (9/5) ^\circ C + 32$

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

• Exact vs. Inexact Numbers: Exact numbers are counted or defined; inexact numbers are measured.

Accuracy, Precision, and Significant Figures

Proper reporting of measurements is crucial for scientific communication.

• Accuracy: Closeness to the true value.

• Precision: Reproducibility of measurements.

• Significant Figures: Digits that carry meaning in a measurement; rules for addition, subtraction, multiplication, and division.

Dimensional Analysis

Dimensional analysis is used to convert units and solve quantitative problems.

• Conversion Factors: Ratios used to express the same quantity in different units.

• Multiple Step Conversions: Chain conversion factors to solve complex problems.

Atomic Theory and Structure

Atomic theory explains the nature and behavior of atoms.

• Dalton's Atomic Theory: Matter is composed of atoms; atoms of each element are identical; atoms combine in simple ratios to form compounds.

• Law of Conservation of Mass: Mass is conserved in chemical reactions.

• Law of Definite Proportions: Compounds have fixed ratios of elements.

• Law of Multiple Proportions: Elements can combine in different ratios to form different compounds.

• Subatomic Particles: Protons (+), neutrons (0), electrons (-); location and properties.

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

• Atomic Number ($Z$): Number of protons.

• Mass Number ($A$): Number of protons plus neutrons.

• Average Atomic Mass: Weighted average based on isotopic abundance.

The Periodic Table

The periodic table organizes elements by atomic number and properties.

• Groups and Periods: Groups are columns; periods are rows.

• Metals, Nonmetals, Metalloids: Location and properties.

• Special Group Names: Alkali metals (Group 1), alkaline earth metals (Group 2), halogens (Group 17), noble gases (Group 18).

• Periodic Trends: Atomic radius, ionization energy, electron affinity, electronegativity.

• Effective Nuclear Charge ($Z_{eff}$): $Z_{eff} = Z - S$ (where $S$ is the shielding constant).

Chemical Bonding

Chemical bonds form between atoms to create molecules and compounds.

• Ionic Bonds: Formed by transfer of electrons between metals and nonmetals.

• Covalent Bonds: Formed by sharing electrons between nonmetals.

• Polar vs. Nonpolar Covalent Bonds: Determined by electronegativity difference.

• Electronegativity: Tendency of an atom to attract electrons in a bond.

• Lewis Dot Structures: Visual representation of valence electrons and bonding.

• Formal Charge: $\text{Formal Charge} = \text{Valence electrons} - (\text{Lone pair electrons} + \frac{1}{2} \text{Bonding electrons})$

• Resonance Structures: Multiple valid Lewis structures for a molecule.

• VSEPR Theory: Predicts molecular geometry based on electron pair repulsion.

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals.

• Bond Strength and Length: Multiple bonds are stronger and shorter than single bonds.

Stoichiometry and Chemical Reactions

Stoichiometry involves quantitative relationships in chemical reactions.

• Balancing Chemical Equations: Ensure the same number of atoms of each element on both sides.

• Types of Reactions: Synthesis, decomposition, single replacement, double replacement, combustion.

• Limiting Reactant: The reactant that is completely consumed first.

• Theoretical Yield: Maximum amount of product possible.

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

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

• Solution Preparation: Calculating concentrations and dilutions.

Electrolytes and Solutions

Solutions can conduct electricity depending on the presence of ions.

• Electrolytes: Substances that dissociate into ions in solution; strong, weak, and nonelectrolytes.

• Precipitation Reactions: Formation of an insoluble product.

• Acid-Base Reactions: Neutralization between acids and bases.

• Redox Reactions: Transfer of electrons; oxidation and reduction.

• Displacement Reactions: Single and double displacement.

• Calculating Molarity: Use mass of solute and volume of solution.

Gases and Gas Laws

Gases have unique properties and behaviors described by gas laws.

• Properties of Gases: Compressibility, expansion, low density.

• Pressure Units: Pascal, bar, mm Hg, atm.

• Standard Atmospheric Pressure: 1 atm = 101.325 kPa = 760 mm Hg.

• 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 P and T)

• Ideal Gas Law: $PV = nRT$

• STP Conditions: 0°C (273.15 K) and 1 atm.

• Density of Gases: $\text{Density} = \frac{PM}{RT}$

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

Quantum Mechanics and Atomic Structure

Quantum mechanics explains the behavior of electrons in atoms.

• Wavelength ($\lambda$) and Frequency ($\nu$): $c = \lambda \nu$

• Electromagnetic Spectrum: Range of all types of electromagnetic radiation.

• Energy of a Photon: $E = h\nu$

• Quantum Numbers: Principal ($n$), angular ($l$), magnetic ($m_l$), spin ($m_s$).

• Electron Configuration: Arrangement of electrons in orbitals; use Aufbau principle, Pauli exclusion, Hund's rule.

• Periodic Trends Explained by Quantum Theory: Atomic radius, ionization energy, electron affinity.

Molecular Structure and Bonding Theories

Molecular geometry and bonding theories predict the shapes and properties of molecules.

• Lewis Dot Structures: Show bonding and lone pairs.

• VSEPR Theory: Predicts shapes based on electron pair repulsion.

• Hybridization: $sp$, $sp^2$, $sp^3$, etc.

• Polarity: Determined by shape and bond dipoles.

• Intermolecular Forces: Hydrogen bonding, dipole-dipole, London dispersion.

• Boiling and Melting Points: Influenced by intermolecular forces.

Additional info:

• This guide covers all major topics listed in a standard General Chemistry curriculum, including measurement, atomic theory, periodic trends, chemical bonding, stoichiometry, solutions, gases, and quantum mechanics.

• Students should be familiar with definitions, formulas, and problem-solving strategies for each topic.