CH101-003 Spring 2026 Final Exam Study Guide: General Chemistry I

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

Significant Figures and Rounding

Accurate scientific measurements require proper use of significant figures and correct rounding procedures.

• 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 to the correct number of significant figures based on the operation performed (addition/subtraction: least decimal places; multiplication/division: least significant figures).

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: $\text{Atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

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 configurations show the arrangement of electrons in atoms and ions.

• Aufbau Principle: Electrons fill lowest energy orbitals first.

• Hund’s Rule: Electrons occupy 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

Compounds are classified as ionic or molecular based on their bonding.

• Ionic Compounds: Formed from metals and nonmetals; consist of ions held by electrostatic forces.

• Molecular Compounds: Formed from nonmetals; consist of molecules held by covalent bonds.

Example: NaCl (ionic), H2O (molecular)

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

Electronegativity and Lewis Structures

Electronegativity is the tendency of an atom to attract electrons. Lewis structures represent the arrangement of atoms and electrons in a molecule.

• Electronegativity Trend: 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 electronegativity differences.

Example: 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: Actual yield divided by theoretical yield, times 100%. $\text{Percent yield} = \frac{\text{actual yield}}{\text{theoretical yield}} \times 100\%$

Chapter 8: Introduction to Solutions and Aqueous Reactions

Molarity Calculations

Molarity expresses concentration as moles of solute per liter of solution.

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

Redox and Precipitation Reactions

• Redox Reactions: Involve transfer of electrons.

• Precipitation Reactions: Formation of insoluble product (precipitate).

Acid/Base Titrations

Titrations determine concentration by reacting a known volume with a solution of known concentration.

Chapter 9: Thermochemistry

Energy, Heat, and Work

Thermochemistry studies energy changes in chemical reactions.

• Internal Energy Change: $\Delta E = q + w$ Sign convention: q (heat) is positive if absorbed, w (work) is positive if done on the system.

• Bomb Calorimeter: Measures $\Delta E$ at constant volume ($q_v$).

• Coffee-Cup Calorimeter: Measures $\Delta H$ at constant pressure ($q_p$).

Enthalpy and Hess’s Law

• Enthalpy of Reaction ($\Delta H$): Heat change at constant pressure.

• Enthalpy of Formation: Enthalpy change for forming 1 mole of compound from elements.

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

• Hess’s Law: The total enthalpy change is the sum of enthalpy changes for individual steps. $\Delta H_{\text{rxn}} = \sum \Delta H_{\text{steps}}$

Chapter 10: Gases

Gas Laws

Gas behavior is described by several laws.

• 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.

• Assumptions: Gas particles are in constant, random motion; collisions are elastic; volume of particles is negligible.

Gas Mixtures and Partial Pressure

• Partial Pressure: Pressure exerted by each gas in a mixture. $P_{\text{total}} = \sum P_i$

Gas Density and Molar Mass

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

Effusion vs. Diffusion

• Effusion: Movement of gas through a small hole.

• Diffusion: Mixing of gases.

• Graham’s Law: $\frac{\text{rate}_1}{\text{rate}_2} = \sqrt{\frac{M_2}{M_1}}$