CH101-003 Spring 2026 Final Exam Study Guide: Comprehensive Review of General Chemistry Topics

Chapter E: Essentials - Units, Measurements, and Problem Solving

Significant Figures and Rounding

Accurate measurement and calculation in chemistry require proper use of significant figures and rounding rules.

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

• Rules:

  • Nonzero digits are always significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

• Rounding: Round the final answer to the correct number of significant figures based on the operation performed.

Example: 0.00450 has three significant figures.

Chapter 1: Atoms

Isotopes and Weighted Atomic Mass

Atoms of the same element can have different numbers of neutrons, resulting in isotopes. The weighted atomic mass reflects the average mass of all isotopes, weighted by their natural abundance.

• Isotope: Atoms with the same number of protons but different numbers of neutrons.

• Weighted Atomic Mass: Calculated as: $\text{Atomic mass} = \sum (\text{fractional abundance} \times \text{isotope mass})$

Example: Carbon-12 and Carbon-13 are isotopes of carbon.

Chapter 2: The Quantum-Mechanical Model of the Atom

Emission Spectra and Bohr Model

The Bohr model explains the emission spectra of hydrogen by quantizing electron energy levels.

• Emission Spectra: Light emitted when electrons transition between energy levels.

• Bohr Model: Electrons orbit the nucleus in fixed energy levels.

• Energy Change: $\Delta E = -R_H \left( \frac{1}{n_f^2} - \frac{1}{n_i^2} \right)$

Quantum Numbers

Quantum numbers describe the properties of atomic orbitals and electrons.

• Principal quantum number (n): Energy level (n = 1, 2, 3, ...)

• Angular momentum quantum number (l): Shape of orbital (l = 0, 1, ..., n-1)

• Magnetic quantum number (ml): Orientation (ml = -l to +l)

• Spin quantum number (ms): Electron spin (ms = +1/2 or -1/2)

Chapter 3: Periodic Properties of the Elements

Trends in Atomic Properties

The periodic table reveals trends in atomic radius, ionization energy, and electron affinity.

• Atomic Radius: Decreases across a period, increases down a group.

• Ionization Energy: Increases across a period, decreases down a group.

• Electron Affinity: Generally becomes more negative across a period.

Electron Configuration

Electron configuration describes the arrangement of electrons in atoms and ions.

• Hund’s Rule: Electrons fill degenerate orbitals singly before pairing.

• Pauli Exclusion Principle: No two electrons in an atom can have the same set of quantum numbers.

Example: The electron configuration of O: 1s2 2s2 2p4

Chapter 4: Molecules and Compounds

Ionic and Molecular Compounds

Chemical compounds are classified as ionic or molecular based on their bonding.

• Ionic Compounds: Formed from metals and nonmetals; consist of ions.

• Molecular Compounds: Formed from nonmetals; consist of molecules.

Example: NaCl (ionic), H2O (molecular)

Chapter 5: Chemical Bonding I - Lewis Structures and Molecular Shapes

Electronegativity and Lewis Structures

Electronegativity determines bond polarity. Lewis structures represent the arrangement of atoms and electrons.

• Electronegativity: Increases across a period, decreases down a group.

• Lewis Structures: Show bonding and lone pairs.

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

• Formal Charge: Calculated as: $\text{Formal charge} = \text{valence electrons} - (\text{nonbonding electrons} + \frac{1}{2}\text{bonding electrons})$

VSEPR and Molecular Shape

VSEPR theory predicts molecular shapes based on electron pair repulsion.

• VSEPR: Valence Shell Electron Pair Repulsion theory.

• Polarity: Determined by shape and bond polarity.

Example: CO2 is linear and nonpolar; H2O is bent and polar.

Chapter 6: Chemical Bonding II - Valence Bond and Molecular Orbital Theory

Hybridization

Hybridization explains the mixing of atomic orbitals to form new hybrid orbitals in molecules.

• Types: sp, sp2, sp3, etc.

• Example: Carbon in methane (CH4) is sp3 hybridized.

Chapter 7: Chemical Reactions and Chemical Quantities

Limiting Reactant, Theoretical Yield, and Percent Yield

Stoichiometry allows calculation of reactant and product quantities.

• Limiting Reactant: The reactant that determines the maximum amount of product.

• Theoretical Yield: Maximum possible product from limiting reactant.

• Percent Yield: Calculated as: $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$

Chapter 8: Introduction to Solutions and Aqueous Reactions

Molarity Calculations

Molarity expresses concentration of a solution.

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

Redox and Precipitation Reactions

• Redox Reaction: Involves transfer of electrons.

• Precipitation Reaction: Formation of an insoluble product.

Acid/Base Titrations

Titrations determine concentration of an acid or base using a reaction with a known solution.

• Endpoint: Point at which reaction is complete.

Chapter 9: Thermochemistry

Energy, Enthalpy, and Calorimetry

Thermochemistry studies energy changes in chemical reactions.

• Internal Energy: $\Delta E = q + w$ (q = heat, w = work)

• Sign Convention: q > 0: heat absorbed; w > 0: work done on system.

• Bomb Calorimeter: $\Delta E = q_v$ (constant volume)

• Coffee-Cup Calorimeter: $\Delta H = q_p$ (constant pressure)

• Enthalpy of Reaction: $\Delta H_{rxn}$

• Enthalpy of Formation: $\Delta H_f$

• Bond Dissociation Energy: Energy required to break a bond.

• Hess’s Law: $\Delta H_{rxn} = \sum \Delta H_{steps}$

Chapter 10: Gases

Gas Laws and Properties

Gas behavior is described by several laws and theories.

• Ideal Gas Law: $PV = nRT$

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Kinetic Molecular Theory: Explains gas properties based on molecular motion.

• Gas Mixtures and Partial Pressure: $P_{total} = \sum P_i$

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

• Molar Mass: $M = \frac{mRT}{PV}$

• Effusion vs. Diffusion: Effusion is movement through a small hole; diffusion is mixing of gases.

Example: Calculate the partial pressure of O2 in air.

Additional info: These notes expand on the brief study guide points, providing definitions, formulas, and examples for each topic listed. This guide covers the foundational concepts required for a general chemistry final exam.