General Chemistry I: Comprehensive Study Notes (Exam Review and Key Concepts)

Introduction to Matter, Energy, and Measurement

Classification of Matter and Properties

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

• Physical properties are characteristics that can be observed without changing the substance's identity (e.g., melting point, density, state of matter).

• Chemical properties describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity with acids).

• Physical changes do not alter the chemical composition, while chemical changes result in new substances.

• Accuracy refers to how close a measurement is to the true value; precision refers to how close repeated measurements are to each other.

Measurement and Significant Figures

• Measurements must be reported with the correct number of significant figures to reflect the precision of the measuring instrument.

• When performing calculations, the result should be rounded to the appropriate number of significant figures based on the operation (addition/subtraction: decimal places; multiplication/division: significant figures).

• Unit conversions are essential for expressing measurements in different units (e.g., converting inches to centimeters, grams to moles).

Example: Measurement Reading

• When reading a ruler, record all certain digits plus one estimated digit.

Atoms, Molecules, and Ions

Atomic Structure and Subatomic Particles

• Atoms consist of a nucleus (protons and neutrons) and electrons in orbitals.

• Protons (positive charge), neutrons (neutral), and electrons (negative charge) are the main subatomic particles.

• The number of protons defines the element (atomic number), while the sum of protons and neutrons gives the mass number.

• Isotopes are atoms of the same element with different numbers of neutrons.

Atomic Models and Experiments

• Rutherford's gold foil experiment showed that atoms have a small, dense, positively charged nucleus.

• Millikan's oil drop experiment determined the charge of the electron.

• The plum pudding model was disproved by Rutherford's findings.

Electron Configuration and Quantum Numbers

• Electrons are arranged in shells and subshells, described by quantum numbers (n, l, ml, ms).

• Electron configurations follow the Aufbau principle, Pauli exclusion principle, and Hund's rule.

• Example: The electron configuration of Ag (noble gas notation): [Kr] 4d10 5s1

Chemical Nomenclature and Formulas

Naming Compounds

• Ionic compounds: Name the cation first, then the anion (e.g., CaCO3 is calcium carbonate).

• Molecular compounds: Use prefixes to indicate the number of atoms (e.g., SF6 is sulfur hexafluoride).

• Acids: Binary acids (e.g., HCl is hydrochloric acid), oxyacids (e.g., H2SO4 is sulfuric acid).

Writing Formulas from Names

• Balance charges to write correct formulas (e.g., iron(III) carbonate is Fe2(CO3)3).

Chemical Reactions and Stoichiometry

Types of Chemical Reactions

• Combination, decomposition, single displacement, double displacement, and combustion reactions.

• Example: Combustion of C2H5OH (ethanol): $C_2H_5OH_{(l)} + 3O_2_{(g)} \rightarrow 2CO_2_{(g)} + 3H_2O_{(l)}$

Balancing Chemical Equations

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

Stoichiometry and Limiting Reactants

• Use molar ratios from balanced equations to calculate amounts of reactants and products.

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

• Theoretical yield is the maximum amount of product possible; percent yield compares actual to theoretical yield: $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$

Reactions in Aqueous Solution

Types of Equations

• Molecular equation: Shows all reactants and products as compounds.

• Complete ionic equation: Shows all strong electrolytes as ions.

• Net ionic equation: Shows only the species that change during the reaction.

Example: Precipitation Reaction

• Na2CO3(aq) + AgNO3(aq) → 2Ag2CO3(s) + 2NaNO3(aq)

Thermochemistry

Heat, Work, and Internal Energy

• First Law of Thermodynamics: Energy cannot be created or destroyed, only transferred.

• Enthalpy change ($\Delta H$) is the heat change at constant pressure.

• Exothermic reactions release heat ($\Delta H < 0$); endothermic reactions absorb heat ($\Delta H > 0$).

Calorimetry

• Used to measure heat changes in chemical reactions.

• Example: Coffee-cup calorimeter for reactions in solution.

Hess's Law

• The enthalpy change for a reaction is the same, regardless of the pathway taken.

Electronic Structure of Atoms

Quantum Numbers and Orbitals

• Principal quantum number (n): energy level

• Angular momentum quantum number (l): shape of orbital

• Magnetic quantum number (ml): orientation

• Spin quantum number (ms): electron spin

Photoelectric Effect and Atomic Spectra

• Energy of a photon: $E = h\nu$

• Wavelength and frequency are related: $c = \lambda \nu$

• Electrons absorb or emit energy when transitioning between energy levels.

Periodic Properties of the Elements

Trends in the Periodic Table

• Atomic radius decreases across a period, increases down a group.

• Ionization energy increases across a period, decreases down a group.

• Electron affinity and electronegativity generally increase across a period.

Classification of Elements

• Metals, nonmetals, metalloids, transition metals, halogens, noble gases.

Basic Concepts of Chemical Bonding

Ionic and Covalent Bonds

• Ionic bonds: Transfer of electrons from metal to nonmetal.

• Covalent bonds: Sharing of electrons between nonmetals.

• Lewis structures represent valence electrons and bonding in molecules.

Bond Polarity and Molecular Shape

• Electronegativity differences determine bond polarity.

• VSEPR theory predicts molecular shapes based on electron pair repulsion.

Molecular Geometry and Bonding Theories

Hybridization and Molecular Orbitals

• Atomic orbitals combine to form hybrid orbitals (e.g., sp3, sp2).

• Molecular orbital theory explains bonding in terms of bonding and antibonding orbitals.

Sigma and Pi Bonds

• Sigma (σ) bonds are single bonds formed by head-on overlap.

• Pi (π) bonds are formed by side-on overlap in double and triple bonds.

Gases

Gas Laws

• Ideal Gas Law: $PV = nRT$

• Other laws: Boyle's Law, Charles's Law, Avogadro's Law.

• Partial pressure: $P_{\text{total}} = \sum P_i$

Kinetic Molecular Theory

• Explains properties of gases in terms of particle motion.

• Root-mean-square speed: $u_{rms} = \sqrt{\frac{3RT}{M}}$

Liquids, Solids, and Intermolecular Forces

Types of Solids

Type of Solid	What are the particles?	Explain what holds the particles together	Notable properties and reasons
Ionic	Cations and anions	Electrostatic attractions	High melting point, brittle, conduct electricity when molten
Metallic	Metal atoms	Delocalized electrons	Malleable, ductile, good conductors
Molecular	Molecules	Intermolecular forces (e.g., hydrogen bonds, dipole-dipole, London dispersion)	Low melting/boiling points, poor conductors
Network covalent	Atoms	Covalent bonds	Very high melting points, hard, poor conductors

Intermolecular Forces

• Types: London dispersion, dipole-dipole, hydrogen bonding.

• Stronger intermolecular forces lead to higher boiling points and lower vapor pressures.

Lab Techniques and Procedures

Measurement and Experimental Error

• Systematic error affects accuracy; random error affects precision.

• Proper use of significant figures and careful measurement reduce error.

Mathematical Operations and Functions

Dimensional Analysis

• Used to convert between units using conversion factors.

• Example: To convert density from g/cm3 to lb/in3, use appropriate conversion factors for mass and volume.

Summary Table: Types of Solids

Type of Solid	Particles	Forces Holding Together	Properties
Ionic	Cations and anions	Electrostatic attraction	High melting point, brittle, conduct when molten
Metallic	Metal atoms	Metallic bonding (sea of electrons)	Malleable, ductile, conduct electricity
Molecular	Molecules	Intermolecular forces	Low melting/boiling points, soft, nonconductors
Network covalent	Atoms	Covalent bonds	Very high melting points, hard, nonconductors

Additional info:

• These notes are based on a comprehensive set of general chemistry exam questions, covering all foundational topics in a typical college-level General Chemistry I course.

• Students should practice applying these concepts to problems involving nomenclature, stoichiometry, atomic structure, periodic trends, bonding, states of matter, and thermochemistry.