General Chemistry I: Final Exam Study Guide (Chapters 1–11)

Chapter 1: Chemical Tools – Experimentation & Measurement

SI Units and Prefixes

The International System of Units (SI) is the standard for scientific measurements. Prefixes are used to denote multiples or fractions of base units.

• Base Units: meter (m), kilogram (kg), second (s), mole (mol), kelvin (K), ampere (A), candela (cd)

• Common Prefixes:

  Prefix	Symbol	Multiplier
  kilo	k	$10^3$
  centi	c	$10^{-2}$
  milli	m	$10^{-3}$
  micro	μ	$10^{-6}$
  nano	n	$10^{-9}$

Density

• Definition: Density is mass per unit volume.

• Formula: $d = \frac{m}{V}$

Significant Figures

• Rules for determining the number of meaningful digits in a measurement.

• Important for reporting calculated results accurately.

Unit Conversions & Dimensional Analysis

• Method for converting between units using conversion factors.

• Ensures consistency and accuracy in calculations.

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

Chapter 2: Atoms, Molecules & Ions

Properties of Matter

• Physical Properties: Characteristics observed without changing composition (e.g., melting point, density).

• Chemical Properties: Characteristics observed during chemical changes (e.g., flammability).

Atomic Structure

• Atoms consist of protons, neutrons, and electrons.

• Atomic Number (Z): Number of protons.

• Mass Number (A): Number of protons + neutrons.

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

Chemical Nomenclature

• Rules for naming compounds and writing chemical formulas.

• Example: NaCl is sodium chloride.

Mole Concept

• 1 mole = $6.022 \times 10^{23}$ entities (Avogadro's number).

• Relates mass, number of particles, and volume for substances.

Chapter 3: Mass Relationships in Chemical Reactions

Balancing Chemical Equations

• Ensures the same number of each atom on both sides of the equation.

• Law of Conservation of Mass: Matter is neither created nor destroyed.

Stoichiometry

• Calculating quantities of reactants and products in chemical reactions.

• Key Steps:

  • Convert mass to moles using molar mass.

  • Use mole ratios from the balanced equation.

  • Convert moles back to mass or volume as needed.

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

• Theoretical Yield: Maximum amount of product possible.

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

Empirical and Molecular Formulas

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

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

Chapter 4: Reactions in Aqueous Solution

Types of Chemical Reactions

• Precipitation Reactions: Formation of an insoluble solid (precipitate) from two solutions.

• Acid-Base Reactions: Transfer of protons (H+) between reactants.

• Oxidation-Reduction (Redox) Reactions: Transfer of electrons between substances.

Solubility Rules

• Guidelines for predicting whether an ionic compound will dissolve in water.

Net Ionic Equations

• Show only the species that actually participate in the reaction.

Identifying Precipitation, Acid-Base, and Redox Reactions

• Use solubility rules, acid-base definitions, and oxidation numbers to classify reactions.

Chapter 5: Periodicity & Electronic Structure of Atoms

Periodic Trends

• Trends in properties such as atomic radius, ionization energy, and electron affinity across periods and groups.

• 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.

Quantum Numbers

• Describe the energy and shape of atomic orbitals.

• Principal (n): Energy level

• Angular Momentum (l): Shape

• Magnetic (ml): Orientation

• Spin (ms): Electron spin direction

Electron Configurations

• Arrangement of electrons in orbitals.

• Follows the Aufbau principle, Pauli exclusion principle, and Hund's rule.

Chapter 6: Ionic Compounds – Periodic Trends and Bonding Theory

Ionic Bonding

• Electrostatic attraction between cations and anions.

• Formation of ionic compounds from metals and nonmetals.

• Lattice Energy: Energy required to separate one mole of an ionic solid into gaseous ions.

Periodic Properties

• Trends in ionic size, charge, and effective nuclear charge.

Chapter 7: Covalent Bonding and Electron-Dot Structures

Covalent Bonding

• Sharing of electron pairs between atoms.

• Bond Length and Strength: Shorter bonds are stronger; multiple bonds (double, triple) are shorter and stronger than single bonds.

Lewis Structures

• Diagrams showing the arrangement of valence electrons among atoms in a molecule.

• Used to predict molecular shape and reactivity.

Resonance

• Some molecules can be represented by two or more valid Lewis structures.

Chapter 8: Covalent Compounds – Bonding Theories and Molecular Structure

VSEPR Theory

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

• Common Shapes: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Bond Polarity and Molecular Polarity

• Polarity depends on differences in electronegativity and molecular geometry.

Intermolecular Forces

• Forces between molecules: dipole-dipole, hydrogen bonding, London dispersion forces.

Chapter 9: Thermochemistry – Chemical Energy

Energy Changes in Chemical Reactions

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

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

Calorimetry

• Measurement of heat flow in a chemical reaction.

• Formula: $q = mc\Delta T$

Hess's Law

• The total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps.

Chapter 10: Gases – Their Properties & Behavior

Gas Laws

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

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

• Ideal Gas Law: $PV = nRT$

• Dalton's Law of Partial Pressures: $P_{\text{total}} = \sum P_i$

Kinetic Molecular Theory

• Explains the behavior of gases in terms of particle motion and energy.

Chapter 11: Liquids & Phase Changes

Properties of Liquids

• Viscosity, surface tension, vapor pressure, boiling point.

Phase Changes

• Energy changes associated with melting, freezing, vaporization, condensation, sublimation, and deposition.

• Calculating Heat of Phase Change: $q = n\Delta H_{\text{phase}}$

Phase Diagrams

• Graphs showing the state of a substance at various temperatures and pressures.

• Interpretation of triple point, critical point, and phase boundaries.