General Chemistry I: Comprehensive Study Guide

Chapter 1 – Introduction: Matter, Energy, and Measurement

Key Concepts in Measurement and Matter

• Matter is anything that has mass and occupies space. It can be classified as elements, compounds, or mixtures (homogeneous or heterogeneous).

• Units of Measurement: Mass (g, kg), Distance (cm, m, km), Volume (L, mL), Temperature (°C, K, °F).

• Pure Substances vs. Mixtures: Pure substances have a fixed composition; mixtures can vary in composition.

• Significant Figures: Important for reporting data accurately after calculations.

• Density: Defined as mass per unit volume. $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Unit Conversions: Use dimensional analysis to convert between units.

Example:

Convert 25.0 cm to meters: $25.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.250\ \text{m}$

Chapter 2 – Atoms, Molecules, and Ions

Atomic Structure and Periodic Table

• Atoms: Consist of protons, neutrons, and electrons. Atomic number = number of protons; mass number = protons + neutrons.

• Ions: Atoms or molecules that have gained or lost electrons. Cations are positively charged; anions are negatively charged.

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

• Periodic Table: Organizes elements by increasing atomic number and similar properties.

• Formulas: Know how to write and name ionic and molecular compounds, including acids.

Example:

Sodium chloride: Na+ and Cl− combine to form NaCl.

Chapter 3 – Chemical Reactions and Reaction Stoichiometry

Balancing Equations and Stoichiometric Calculations

• Balancing Chemical Equations: Ensure the same number of each atom on both sides of the equation.

• Types of Reactions: Decomposition, replacement, combination, etc.

• Avogadro’s Number: $6.022 \times 10^{23}$ particles/mol.

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

• Empirical and Molecular Formulas: Empirical is the simplest ratio; molecular is the actual number of atoms.

• Stoichiometry: Use balanced equations to calculate reactant and product quantities.

Key Equations:

• $\text{number of moles} = \frac{\text{mass}}{\text{molar mass}}$

• $\text{Molarity (M)} = \frac{\text{moles of solute}}{\text{volume of solution in L}}$

Example:

How many moles in 18 g of H2O? $\frac{18\ \text{g}}{18.0\ \text{g/mol}} = 1.00\ \text{mol}$

Chapter 4 – Reactions in Aqueous Solution

Precipitation, Acid-Base, and Redox Reactions

• Solubility Rules: Used to predict whether a precipitate will form in a reaction.

• Types of Reactions: Precipitation, acid-base (neutralization), and redox reactions.

• Net Ionic Equations: Show only the species that change during the reaction.

• Concentration Calculations: Use molarity to determine amounts of reactants and products.

Example:

Mixing NaCl and AgNO3 forms AgCl(s) precipitate: $\text{Ag}^+ (aq) + \text{Cl}^- (aq) \rightarrow \text{AgCl} (s)$

Chapter 6 – Electronic Structure of Atoms

Quantum Theory and Electron Configuration

• Electromagnetic Radiation: Light has both wave and particle properties.

• Key Equations:

  • $c = \lambda \nu$ (speed of light = wavelength × frequency)

  • $E = h\nu$ (energy of a photon)

  • $\lambda = \frac{h}{mv}$ (de Broglie wavelength)

• Quantum Numbers: Describe the energy, shape, and orientation of orbitals.

• Electron Configuration: Follows the Aufbau principle, Pauli exclusion principle, and Hund’s rule.

Example:

Electron configuration of O: 1s2 2s2 2p4

Chapter 7 – Periodic Properties of the Elements

Trends in the Periodic Table

• Effective Nuclear Charge (Zeff): The net positive charge experienced by valence electrons.

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

• Ionization Energy: Energy required to remove an electron; increases across a period, decreases down a group.

• Electron Affinity: Energy change when an electron is added to an atom.

• Metallic Character: Increases down a group, decreases across a period.

Example:

Na has a larger atomic radius than Cl; Cl has a higher ionization energy than Na.

Chapter 8 – Basic Concepts of Chemical Bonding

Ionic and Covalent Bonding

• Ionic Bonds: Formed by transfer of electrons from metals to nonmetals.

• Covalent Bonds: Formed by sharing electrons between nonmetals.

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

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

• Bond Polarity: Difference in electronegativity determines if a bond is polar or nonpolar.

• Lewis Structures: Show arrangement of electrons in molecules; formal charges help determine the most stable structure.

Example:

CO2 is a nonpolar molecule with polar bonds.

Chapter 9 – Molecular Geometry and Bonding Theories

VSEPR Theory and Molecular Shapes

• VSEPR Theory: Predicts the shape of molecules based on electron pair repulsion.

• Common Geometries:

  Electron Groups	Lone Pairs	Geometry	Bond Angle	Example
  2	0	Linear	180°	CO2
  3	0	Trigonal planar	120°	BF3
  4	0	Tetrahedral	109.5°	CH4
  4	1	Trigonal pyramidal	107°	NH3
  4	2	Bent	104.5°	H2O

• Hybridization: Atomic orbitals mix to form new hybrid orbitals (e.g., sp, sp2, sp3).

• Molecular Polarity: Determined by shape and bond polarity.

Example:

CH4 is tetrahedral and nonpolar; H2O is bent and polar.

Chapter 10 – Gases

Gas Laws and Kinetic Molecular Theory

• Gas Laws:

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

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

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

  • Ideal Gas Law: $PV = nRT$

• Partial Pressure: $P_A = X_A P_{\text{total}}$ where $X_A$ is the mole fraction.

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

• Kinetic Molecular Theory: Explains the behavior of gases in terms of particle motion.

Example:

Calculate the pressure exerted by 2.0 mol of gas in a 5.0 L container at 300 K: $P = \frac{nRT}{V}$

Chapter 11 – Liquids and Intermolecular Forces

Types of Intermolecular Forces and Properties of Liquids

• Intermolecular Forces:

  Dipole-dipole attractions	Dispersion forces
  Attractive forces between molecules with permanent dipole moments (polar molecules)	Attractive forces between all molecules due to instantaneous dipoles
  Generally stronger	Generally weaker
  Permanent dipoles must exist	Instantaneous dipoles must exist

• Hydrogen Bonding: Special dipole-dipole interaction involving H bonded to N, O, or F.

• Factors Affecting Boiling and Melting Points: Stronger intermolecular forces lead to higher boiling/melting points.

• Phase Changes: Melting, freezing, vaporization, condensation, sublimation, deposition.

• Phase Diagrams: Show the state of a substance at various temperatures and pressures; critical point, triple point, etc.

Example:

Water has a high boiling point due to hydrogen bonding.

Key Equations to Remember

• $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• $\text{Moles} = \frac{\text{mass}}{\text{molar mass}}$

• $\text{Molarity (M)} = \frac{\text{moles of solute}}{\text{volume of solution in L}}$

• $PV = nRT$

• $E = h\nu$

• $c = \lambda \nu$

• $\lambda = \frac{h}{mv}$

Additional info:

• This study guide covers the foundational topics for a first-semester General Chemistry course, including measurement, atomic structure, chemical reactions, bonding, molecular geometry, gases, and intermolecular forces.

• Tables and equations have been recreated and expanded for clarity and completeness.