General Chemistry I: Comprehensive Study Notes (Chapters 1–11)

CHAPTER 1 — Matter, Measurements, and Problem Solving

Scientific Method and Measurement

The scientific method is a systematic approach to understanding the natural world through observation and experimentation. Accurate measurement and data analysis are foundational to chemistry.

• Scientific Method Steps: Observation → Hypothesis → Experiment → Analyze Results → Conclusion

• States of Matter:

  • Solid: Fixed shape and volume

  • Liquid: Fixed volume, variable shape

  • Gas: Variable shape and volume

• Physical vs. Chemical Changes:

  • Physical: Melting, boiling, dissolving

  • Chemical: Combustion, rusting, gas/precipitate formation

• Intensive Properties: Independent of amount (e.g., density, boiling point, color)

• Extensive Properties: Dependent on amount (e.g., mass, volume, length)

• SI Units & Metric Prefixes: Standardized units for scientific measurement (e.g., meter, kilogram, second)

• Dimensional Analysis: Systematic unit conversion; always cancel units step-by-step.

• Significant Figures:

  • Nonzero digits are significant

  • Interior zeros are significant

  • Leading zeros are not significant

  • Trailing zeros after a decimal are significant

• Accuracy vs. Precision: Accuracy is closeness to true value; precision is reproducibility.

Example: Converting 25.0°C to Kelvin: $T(K) = 25.0 + 273.15 = 298.15$

CHAPTER 2 — Atoms, Molecules, and Ions

Atomic Structure and Chemical Nomenclature

Understanding the structure of atoms and how they combine to form compounds is fundamental to chemistry.

• Dalton’s Atomic Theory: Atoms are indivisible particles; atoms of the same element are identical.

• Subatomic Particles: Proton (+), Neutron (0), Electron (−)

• Atomic Number (Z): Number of protons

• Mass Number (A): Protons + Neutrons

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

• Average Atomic Mass: Weighted average based on isotope abundance

• Ionic Compounds: Metal + nonmetal; electrons transferred

• Molecular Compounds: Nonmetal + nonmetal; electrons shared; use prefixes (mono-, di-, tri-, etc.)

• Acid Naming:

  • Binary: hydro___ic acid

  • Oxyacid: -ate → -ic acid, -ite → -ous acid

Example: $^{12}\mathrm{C}$ and $^{13}\mathrm{C}$ are isotopes of carbon.

CHAPTER 3 — Stoichiometry

Quantitative Chemical Relationships

Stoichiometry involves calculations based on balanced chemical equations, relating masses, moles, and numbers of particles.

• Balancing Equations: Conserve atoms; never change subscripts.

• Mole Roadmap: mass ↔ moles ↔ particles (using Avogadro’s number: $6.022 \times 10^{23}$)

• Empirical Formula Steps:

  1. Convert % to grams

  2. Convert grams to moles

  3. Divide by smallest number of moles

  4. Multiply to get whole numbers

• Limiting Reactant: The reactant that produces the least amount of product

• Theoretical Yield: Maximum possible product

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

Example: If 10.0 g of A reacts with 10.0 g of B, calculate which is limiting and the percent yield if 8.0 g of product is obtained.

CHAPTER 4 — Reactions in Aqueous Solution

Solution Chemistry and Ionic Reactions

Many chemical reactions occur in water. Understanding solubility, electrolytes, and ionic equations is essential.

• Electrolytes:

  • Strong: Strong acids, strong bases, soluble ionic compounds

  • Weak: Weak acids and bases

  • Nonelectrolytes: Sugars, alcohols

• Solubility Rules: (see table below)

• Net Ionic Equations: Show only species that change during the reaction; cancel spectator ions.

• Oxidation Numbers: Assign based on rules (e.g., free element = 0, O = -2, H = +1, Group 1 = +1, Group 2 = +2)

• Redox Reactions: Oxidation = loss of electrons; Reduction = gain of electrons (OIL RIG)

• Molarity: $M = \frac{\text{mol solute}}{\text{L solution}}$

• Dilution Formula: $M_1V_1 = M_2V_2$

• Titration: At equivalence, moles acid = moles base; indicators signal endpoint

Solubility Rule	Exceptions
Most nitrates (NO3-) and acetates (C2H3O2-)	None
Most Group 1A and NH4+ compounds	None
Most chlorides, bromides, iodides	Ag+, Hg22+, Pb2+
Most sulfates	Sr2+, Ba2+, Pb2+, Hg22+
Most hydroxides	Except Ca2+, Sr2+, Ba2+
Most carbonates, phosphates, sulfides, chromates, dichromates, sulfites	Insoluble

CHAPTER 5 — Thermochemistry

Energy Changes in Chemical Reactions

Thermochemistry studies heat and energy changes during chemical reactions.

• Exothermic: Releases heat ($\Delta H$ negative)

• Endothermic: Absorbs heat ($\Delta H$ positive)

• Specific Heat (c): Heat required to raise 1 g by 1°C

• Heat Equation: $q = mC\Delta T$

• Calorimetry: $q_{\text{rxn}} = -q_{\text{soln}}$

• Enthalpy from Moles: $q = \text{mol} \times \Delta H$

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

• Bond Enthalpy: $\Delta H = \sum \text{Bonds Broken} - \sum \text{Bonds Formed}$

• Internal Energy: $\Delta E = q + w$

Example: Calculate the heat absorbed when 50.0 g of water is heated from 25°C to 75°C. ($c = 4.18$ J/g·°C)

CHAPTER 6 — Electronic Structure of Atoms

Quantum Mechanics and Electron Configuration

Quantum theory explains the arrangement of electrons in atoms and their energy levels.

• Quantum Numbers:

  • n: Principal quantum number (energy level)

  • l: Angular momentum (0 = s, 1 = p, 2 = d, 3 = f)

  • ml: Magnetic quantum number (orbital orientation)

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

• Orbital Shapes: s = spherical, p = dumbbell

• Electron Configuration Rules:

  • Aufbau Principle: Fill lowest energy orbitals first

  • Hund’s Rule: Fill singly before pairing

  • Pauli Exclusion: Max 2 electrons per orbital, opposite spins

• Photon Energy: $E = h\nu$

• Speed of Light: $c = \nu\lambda$

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

CHAPTER 7 — Periodic Properties of the Elements

Trends in the Periodic Table

Periodic trends help predict element properties based on their position in the periodic table.

• Atomic Radius: Increases down a group, decreases across a period

• Ionization Energy: Increases across, decreases down

• Electron Affinity: Generally more negative across a period

• Metallic Character: Increases down, increases to the left

• Important Exception: Nitrogen has higher ionization energy than oxygen due to half-filled p orbitals

Trend	Across Period	Down Group
Atomic Radius	Decreases	Increases
Ionization Energy	Increases	Decreases
Electron Affinity	More Negative	Less Negative
Metallic Character	Decreases	Increases

CHAPTER 8 — Basic Concepts of Chemical Bonding

Lewis Structures, Resonance, and Bonding

Bonding theories explain how atoms combine and the shapes of molecules.

• Lewis Structure Steps:

  1. Count valence electrons

  2. Choose central atom

  3. Make single bonds

  4. Complete octets

  5. Add multiple bonds if needed

• Resonance: Multiple valid Lewis structures; actual structure is a hybrid

• Formal Charge: $\text{FC} = \text{valence e}^- - \text{assigned e}^-$

• Bond Trends: Shorter bond = stronger bond; triple > double > single

• Octet Rule Exceptions: H (2e-), Be (4e-), B (6e-), expanded octets for period 3+

CHAPTER 9 — Molecular Geometry and Bonding Theories

VSEPR, Hybridization, and Molecular Polarity

Molecular geometry determines the shape and polarity of molecules, affecting their physical and chemical properties.

• VSEPR Shapes:

  • 2 groups: Linear (180°)

  • 3 groups: Trigonal planar (120°)

  • 4 groups: Tetrahedral (109.5°)

  • 5 groups: Trigonal bipyramidal

  • 6 groups: Octahedral

• Common Shapes: Bent, trigonal pyramidal, T-shaped, seesaw

• Hybridization: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral)

• Molecular Polarity: Depends on bond polarity and geometry

• Molecular Orbitals: Bonding orbitals stabilize; antibonding destabilize

• Paramagnetic: Unpaired electrons; Diamagnetic: All electrons paired

Electron Groups	Electron Geometry	Molecular Geometry
2	Linear	Linear
3	Trigonal Planar	Trigonal Planar / Bent
4	Tetrahedral	Tetrahedral / Trigonal Pyramidal / Bent
5	Trigonal Bipyramidal	Seesaw / T-shaped / Linear
6	Octahedral	Square Pyramidal / Square Planar

CHAPTER 10 — Gases

Gas Laws and Kinetic Theory

The behavior of gases is described by several empirical laws and the kinetic molecular theory.

• Boyle’s Law: $P_1V_1 = P_2V_2$ (inverse relationship)

• Charles’s Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (direct relationship)

• Avogadro’s Law: $\frac{n_1}{V_1} = \frac{n_2}{V_2}$

• Ideal Gas Law: $PV = nRT$

• Dalton’s Law: $P_{\text{total}} = P_1 + P_2 + ...$

• Kinetic Molecular Theory: Gas particles are tiny, in constant random motion, with elastic collisions and negligible attractions

• Root Mean Square Speed: $\mu = \sqrt{\frac{3RT}{M}}$

• Graham’s Law: Lighter gases diffuse faster

CHAPTER 11 — Liquids, Solids, and Intermolecular Forces

Phases of Matter and Intermolecular Forces

Intermolecular forces determine the physical properties of substances, such as boiling point and solubility.

• Intermolecular Forces (IMF):

  • London dispersion (weakest)

  • Dipole-dipole

  • Hydrogen bonding (requires H bonded to N, O, or F)

  • Ion-dipole (strongest)

• Trends: Stronger IMF → higher boiling point, viscosity, surface tension

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

• Critical Point: No distinction between liquid/gas above this temperature and pressure

• Triple Point: All three phases coexist

IMF	Relative Strength
London Dispersion	Weakest
Dipole-Dipole	Medium
Hydrogen Bonding	Strong
Ion-Dipole	Strongest

Lab Techniques and Procedures

Common Laboratory Equipment and Errors

Familiarity with laboratory equipment and error analysis is essential for accurate experimental work.

• Volumetric Pipet: Highest precision for volume delivery

• Buret: Used for titrations

• Calorimeter: Measures heat changes

• Desiccator: Keeps samples dry

• Common Errors:

  • Overshooting endpoint: Makes concentration too high

  • Wet precipitate: Increases measured mass

  • Contamination: Usually increases error

  • Spilled sample: Lowers calculated amount

Mathematical Operations and Functions

Key Equations and Conversions

• Density: $d = \frac{m}{V}$

• Molarity: $M = \frac{\text{mol solute}}{\text{L solution}}$

• Dilution: $M_1V_1 = M_2V_2$

• Temperature Conversions:

  • Celsius to Kelvin: $T(K) = T(\degree C) + 273.15$

  • Fahrenheit to Celsius: $T(\degree C) = \frac{T(\degree F) - 32}{1.8}$

Final Exam Strategies and High-Yield Topics

• Memorize solubility rules, strong acids/bases, polyatomic ions, gas law equations, thermochemistry equations, Lewis structures, VSEPR shapes, periodic trends, oxidation number rules, and stoichiometry roadmap.

• Common traps: Forgetting mole ratios, using grams instead of moles, not identifying limiting reactant, wrong R value, pressure unit mismatch, incorrect electron count in Lewis structures, sign errors in $\Delta H$, and unit conversion mistakes.

• Test-taking tips: Do easiest questions first, double-check units and signs, eliminate impossible answers, and watch for limiting reactants and total volume changes.

Additional info:

• Mnemonic for diatomic elements: "Have No Fear Of Ice Cold Beer" (H2, N2, F2, O2, I2, Cl2, Br2)

• Visible hydrogen emission lines correspond to the Balmer series.

• At STP (Standard Temperature and Pressure), 1 mol of gas occupies 22.4 L.