General Chemistry Final Exam Review: Comprehensive Study Notes

Chapter 1: Introduction to Chemistry and Matter

Matter and Its Classification

Chemistry is the study of matter, its properties, and the changes it undergoes. Matter can be classified based on its composition and physical state.

• Pure Substances: Have a fixed composition and distinct properties.

  • Elements: Consist of only one type of atom.

    • Monatomic: Elements that exist as single atoms (e.g., He).

    • Diatomic: Elements that exist as molecules of two atoms (e.g., O2).

  • Compounds: Substances composed of two or more elements chemically combined.

    • Binary: Contain two different elements (e.g., NaCl).

    • Ternary: Contain three different elements (e.g., H2SO4).

• Mixtures: Physical combinations of two or more substances.

  • Homogeneous: Uniform composition throughout (e.g., saltwater).

  • Heterogeneous: Non-uniform composition (e.g., salad).

  • Colloids and Suspensions: Mixtures with particles dispersed in another substance; colloids have intermediate particle size, suspensions have larger particles that may settle.

Properties and Changes of Matter

• Physical Properties: Can be observed without changing the substance's identity (e.g., melting point, density).

• Chemical Properties: Describe a substance's ability to undergo chemical changes (e.g., flammability).

• Physical Change: Alters appearance but not composition (e.g., melting ice).

• Chemical Change: Produces new substances (e.g., rusting iron).

States of Matter

• Solid: Fixed shape and volume; particles vibrate in place.

• Liquid: Fixed volume, no fixed shape; particles move more freely.

• Gas: No fixed shape or volume; particles move rapidly and are far apart.

State	Fixed Shape?	Fixed Volume?	Particle Motion
Solid	Yes	Yes	Vibrate in place
Liquid	No	Yes	Move past each other
Gas	No	No	Rapid, random motion

Measuring Matter

• Qualitative: Descriptions without numbers (e.g., color).

• Quantitative: Measurements with numbers (e.g., mass, volume).

• SI System: Standard units for scientific measurements (e.g., kilogram for mass, kelvin for temperature).

• Temperature Scales:

  • Celsius (°C)

  • Fahrenheit (°F)

  • Kelvin (K)

• Conversion Factors: Used to convert between units.

• Significant Figures: Indicate the precision of a measurement.

  • Multiplying/Dividing: Result has as many significant figures as the measurement with the fewest.

  • Adding/Subtracting: Result has as many decimal places as the measurement with the fewest.

• Scientific Notation: Expresses numbers as a product of a coefficient and a power of ten.

• Density as a Conversion Factor: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

Chapter 2: Atoms, Elements, and the Periodic Table

Structure of the Atom

• Subatomic Particles:

  • Protons: Positive charge, located in nucleus, mass ≈ 1 amu.

  • Neutrons: No charge, located in nucleus, mass ≈ 1 amu.

  • Electrons: Negative charge, located outside nucleus, mass ≈ 0.0005 amu.

• Element Symbol: One- or two-letter abbreviation (e.g., H, He).

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Total number of protons and neutrons.

• Neutron Number: $\text{Neutrons} = A - Z$

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

• Average Atomic Mass: Weighted average of isotopes (in amu).

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

• Ground State: Lowest energy state of an atom.

• Ions: Atoms with a net charge due to loss or gain of electrons.

  • Cation: Positively charged (loss of electrons).

  • Anion: Negatively charged (gain of electrons).

The Periodic Table

• Rows/Periods: Horizontal rows; indicate energy levels.

• Columns/Groups/Families: Vertical columns; elements with similar properties.

• Metals, Nonmetals, Metalloids: Classified by physical and chemical properties.

• Hydrogen: Unique, placed above Group 1 but not a metal.

• Groups:

  • Alkali Metals (Group 1): 1 valence electron, +1 charge.

  • Alkaline Earth Metals (Group 2): 2 valence electrons, +2 charge.

  • Halogens (Group 17): 7 valence electrons, -1 charge.

  • Noble Gases (Group 18): 8 valence electrons (except He), inert.

  • Transition Metals: Variable charges.

  • Lanthanides and Actinides: Inner transition metals.

The Mole and Chemical Calculations

• Mole: Amount of substance containing $6.022 \times 10^{23}$ entities (Avogadro's number).

• Conversions:

  • Atoms ↔ Moles: $\text{Moles} = \frac{\text{atoms}}{6.022 \times 10^{23}}$

  • Grams ↔ Moles: $\text{Moles} = \frac{\text{grams}}{\text{molar mass}}$

  • Grams ↔ Atoms: Two-step conversion via moles.

History of Atomic Theory

• Law of Conservation of Mass (Lavoisier): Mass is neither created nor destroyed in chemical reactions.

• Law of Constant Composition: A given compound always contains the same proportion of elements by mass.

Chapter 3: Compounds and Chemical Nomenclature

Chemical Formulas and Molar Mass

• Subscripts: Indicate the number of each atom in a molecule.

• Molar Mass: Mass of one mole of a compound (g/mol).

• Conversions:

  • Grams ↔ Moles ↔ Molecules ↔ Atoms (using Avogadro's number and molar mass).

• Percent Composition: $\%\text{ composition} = \frac{\text{mass of element in 1 mol}}{\text{molar mass}} \times 100\%$

Types of Compounds and Nomenclature

• Ionic Compounds: Formed from metals and nonmetals; transfer of electrons.

• Covalent Compounds: Formed from nonmetals; sharing of electrons.

• Octet Rule: Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons.

• Typical Charges: Groups 1A-7A have predictable charges based on group number.

• Crossover Rule: Used to write formulas for ionic compounds by balancing charges.

• Polyatomic Ions: Use parentheses when more than one is present in a formula.

• Naming Compounds:

  • Covalent: Use prefixes (mono-, di-, tri-, etc.).

  • Ionic Type 1: Metal + nonmetal; use crossover rule.

  • Ionic Type 2: Transition metal + nonmetal; indicate metal's charge with Roman numerals.

  • Ionic Type 3: Compounds with polyatomic ions.

Chapter 4: Chemical Reactions and Stoichiometry

Balancing and Writing Chemical Equations

• Counting Atoms: Ensure the same number of each atom on both sides of the equation.

• Balancing Equations: Adjust coefficients to obey the law of conservation of mass.

• Writing Formulas from Names: Use nomenclature rules to write correct formulas.

Stoichiometry

• Stoichiometry: Quantitative relationships in chemical reactions.

• Mole-to-Mole Conversions: Use coefficients from balanced equations.

• Grams-to-Grams Conversions: Convert grams to moles, use mole ratio, then convert back to grams.

• Limiting Reactant: The reactant that produces the least amount of product; determines the maximum yield.

• Excess Reactant: The reactant left over after the reaction.

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

• Types of Reactions: Synthesis, decomposition, single replacement, double replacement, combustion (not on final).

Chapter 5: Solutions and Acid-Base Chemistry

Molarity and Dilutions

• Molarity (M): $M = \frac{n}{V}$, where n = moles of solute, V = volume in liters.

• Dilute vs. Concentrated: Refers to the amount of solute in solution.

• Dilution Equation: $M_1V_1 = M_2V_2$

Precipitation and Acid-Base Reactions

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

• Solubility Guidelines: Used to predict if a precipitate will form.

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

• Acid-Base Definitions:

  • Arrhenius: Acids produce H+, bases produce OH- in water.

  • Bronsted-Lowry: Acids donate protons, bases accept protons.

• Strong vs. Weak Acids/Bases: Strong acids/bases dissociate completely; weak only partially.

• Water Ionization Constant: $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$

• pH Calculation: $pH = -\log[H^+]$

• Neutralization: Acid + base → salt + water.

Chapter 6: Redox Reactions and Gases

Redox Reactions

• Oxidation: Loss of electrons (OIL: Oxidation Is Losing).

• Reduction: Gain of electrons (RIG: Reduction Is Gaining).

• Identifying Redox:

  • Change in oxidation number.

  • Gain/loss of oxygen.

  • Half-reactions: Electrons on right (oxidation), on left (reduction).

Gases and Gas Laws

• Unit Conversions:

  • Temperature: $T(\text{K}) = T(\degree C) + 273.15$

  • Volume: 1 L = 1000 mL

  • Pressure: 1 atm = 760 Torr = 760 mmHg = 101.325 kPa

• Gas Laws:

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

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

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

  • Gay-Lussac's Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ (at constant V, n)

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

  • Ideal Gas Law: $PV = nRT$

• Law of Combining Volumes: 1 mol of any gas at STP occupies 22.7 L.

• Dalton's Law of Partial Pressures: $P_{total} = P_1 + P_2 + ...$

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

Chapter 7: Thermochemistry

Systems and Energy

• System Types:

  • Open: Exchanges matter and energy with surroundings.

  • Closed: Exchanges energy but not matter.

  • Isolated: Exchanges neither matter nor energy.

• Law of Conservation of Matter: Matter is not created or destroyed.

• First Law of Thermodynamics: Energy is conserved.

• Internal Energy (U): Total energy of a system.

• Work (w): Energy transfer due to force acting over distance.

• Heat (q): Energy transfer due to temperature difference.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

Heat and Calorimetry

• Calorie: 1 cal = 4.18 J

• Heat Equation: $q = mc\Delta T$

  • q = heat (J)

  • m = mass (g)

  • c = specific heat (J/g·K)

  • $\Delta T$ = change in temperature (K or °C)

• Law of Conservation of Energy: $q_{system} = -q_{surroundings}$

• Exothermic: Releases heat (q negative).

• Endothermic: Absorbs heat (q positive).

• Work: Positive when done on the system, negative when done by the system.

Chapter 8: Quantum Mechanics and Atomic Structure

Bohr Model and Electron Configuration

• Bohr Model: Electrons orbit nucleus in quantized energy levels.

• Bohr Notation: Shows electron configuration for atoms and ions.

• Energy Levels: n = 1 to 7; each can hold $2n^2$ electrons.

• Excited State: Electron occupies higher energy level than ground state.

Light and Electromagnetic Spectrum

• Wave Properties: Wavelength (λ), frequency (ν), speed (c).

• Relationship: $c = \lambda \nu$, where $c = 3.00 \times 10^8$ m/s.

• Photon Energy: $E = h\nu$, where $h = 6.626 \times 10^{-34}$ Js.

• Visible Light: Part of the electromagnetic spectrum detectable by the human eye.

• Emission Spectrum: Unique set of wavelengths emitted by an element.

• Flame Test: Identifies elements by color of flame.

• Wave-Particle Duality: Light exhibits both wave and particle properties.

• Heisenberg Uncertainty Principle: Impossible to know both position and momentum of an electron precisely.

• Schrodinger Equation: Describes electron behavior as a wave.

Quantum Numbers and Orbitals

• Principal Quantum Number (n): Energy level (1-7).

• Angular Momentum Quantum Number (l): Sublevel (s=0, p=1, d=2, f=3).

• Magnetic Quantum Number (ml): Orientation of orbital.

• Spin Quantum Number (ms): +1/2 (up) or -1/2 (down).

• Orbital Types:

  • s: 1 orientation

  • p: 3 orientations

  • d: 5 orientations

  • f: 7 orientations

• Aufbau Principle: Electrons fill lowest energy orbitals first.

• Hund's Rule: Electrons occupy degenerate orbitals singly before pairing.

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

• Electron Configuration: Arrangement of electrons in orbitals.

• Orbital Notation: Uses arrows to represent electron spins.

Chapter 9: Periodic Properties of the Elements

Periodic Law and Trends

• Periodic Law: Properties of elements repeat periodically when arranged by atomic number.

• Modern Periodic Table: Arranged by increasing atomic number.

• Trends:

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

  • Effective Nuclear Charge (Z*): Net positive charge experienced by valence electrons.

  • Valence Electrons: Electrons in the outermost shell.

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

  • Electron Affinity/Electronegativity: Tendency to attract electrons; increases across a period, decreases down a group.

  • Magnetic Properties:

    • Diamagnetic: All electrons paired.

    • Paramagnetic: Unpaired electrons present.

Chapter 10: Bonding and Molecular Structure

Lewis Structures and Molecular Geometry

• Octet Rule: Atoms tend to form bonds to achieve 8 valence electrons.

• Lewis Dot Structures: Represent valence electrons as dots around element symbols.

• Multiple Bonds: Double or triple bonds between atoms.

• Polyatomic Ions: Charged species composed of multiple atoms.

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

• Exceptions to Octet Rule: Some atoms have fewer or more than 8 electrons.

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

VSEPR Theory and Molecular Shapes

• VSEPR (Valence Shell Electron Pair Repulsion) Theory: Predicts molecular geometry based on repulsion between electron pairs.

• Common Geometries:

  • Linear: 180° bond angle

  • Tetrahedral: 109.5° bond angle

  • Trigonal Planar: 120° bond angle

  • Trigonal Pyramidal: slightly less than 109.5°

  • Bent: less than 120° or 109.5°, depending on lone pairs