General Chemistry I: Comprehensive Study Guide by Unit

Unit 1: Atoms, Ions, and Chemical Compounds

Atomic Structure and Particles

Atoms are composed of three main subatomic particles: protons, neutrons, and electrons. The number and type of each particle determine the identity and properties of an atom or ion.

• Protons (p+): Positively charged particles found in the nucleus; the atomic number (Z) equals the number of protons.

• Neutrons (n0): Neutral particles in the nucleus; the number of neutrons = mass number (A) - atomic number (Z).

• Electrons (e-): Negatively charged particles in orbitals around the nucleus; in a neutral atom, electrons = protons. In ions, electrons differ by the charge.

• Cations: Positively charged ions (fewer electrons than protons).

• Anions: Negatively charged ions (more electrons than protons).

Example: For 23Na+: 11 protons, 12 neutrons, 10 electrons.

Naming and Writing Formulas for Compounds

• Binary Compounds: Contain two elements (e.g., NaCl, CO2).

• Ternary Compounds: Contain three elements, often with polyatomic ions (e.g., NaNO3, CaCO3).

• Naming: Use systematic rules for ionic and molecular compounds.

Example: K2SO4 is potassium sulfate.

Molar Mass and Conversions

• Molar Mass (g/mol): The mass of one mole of a substance.

• Conversions: Use Avogadro's number ($6.022 \times 10^{23}$ particles/mol) and molar mass for calculations.

Key Equations:

• $\text{moles} = \frac{\text{mass (g)}}{\text{molar mass (g/mol)}}$

• $\text{number of particles} = \text{moles} \times 6.022 \times 10^{23}$

Percent Composition

• Percent Composition: The percent by mass of each element in a compound.

• $\%\,\text{element} = \frac{\text{mass of element in 1 mol}}{\text{molar mass of compound}} \times 100\%$

Example: In H2O, %H = 11.2%, %O = 88.8%.

Unit 2: Electron Configuration and Periodic Properties

Electron Configurations

Electron configurations describe the arrangement of electrons in an atom's orbitals.

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

• Example: Oxygen (Z=8): 1s2 2s2 2p4

Periodic Properties

• Atomic Size: Increases down a group, decreases across a period.

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

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

Application: These trends explain reactivity and bonding behavior.

Energy, Wavelength, and Frequency

• Energy ($E$), wavelength ($\lambda$), and frequency ($\nu$) are related for electromagnetic radiation.

• $E = h\nu$

• $c = \lambda \nu$

• Where $h$ is Planck's constant ($6.626 \times 10^{-34}$ J·s), $c$ is the speed of light ($3.00 \times 10^8$ m/s).

Unit 3: Chemical Reactions and Equations

Writing and Balancing Equations

• Write correct formulas for reactants and products.

• Balance equations to conserve mass.

Types of Reactions: Synthesis, decomposition, single replacement, double replacement, combustion.

Activity Series and Solubility Rules

• Activity Series: Predicts if a metal will displace another in a reaction.

• Solubility Rules: Determine if a compound is soluble in water.

Unit 4: Stoichiometry and Molarity

Stoichiometry Calculations

• Use balanced equations to relate moles of reactants and products.

• Convert between grams, moles, and particles as needed.

Molarity

• Molarity (M): $M = \frac{\text{moles of solute}}{\text{liters of solution}}$

• Used in solution stoichiometry calculations.

Unit 5: Thermochemistry and Reaction Spontaneity

Enthalpy Changes ($\Delta H$)

• Use Hess's Law, heats of formation, or calorimetry to determine $\Delta H$.

• Hess's Law: The total enthalpy change is the sum of individual steps.

• Calorimetry: $q = mc\Delta T$

Exothermic vs. Endothermic Reactions

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

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

Entropy and Spontaneity

• Entropy ($\Delta S$): Measure of disorder; increases with more microstates.

• Gibbs Free Energy: $\Delta G = \Delta H - T\Delta S$

• Spontaneous if $\Delta G < 0$.

Unit 6: Chemical Bonding and Molecular Structure

Lewis Structures and Molecular Geometry

• Draw Lewis dot structures for molecules and ions.

• Use VSEPR theory to predict geometry (e.g., linear, bent, tetrahedral).

Types of Bonding

• Ionic: Transfer of electrons (metal + nonmetal).

• Covalent: Sharing of electrons (nonmetal + nonmetal).

• Electronegativity differences help distinguish bond type.

Molecular Polarity and Intermolecular Forces

• Determine net dipole moment from structure.

• Types of intermolecular forces: London dispersion, dipole-dipole, hydrogen bonding.

• Polarity affects solubility in water (polar) or hexane (nonpolar).

Unit 7: Properties of Liquids, Solids, and Gases

Vapor Pressure and Properties of Liquids

• Vapor pressure: Pressure exerted by a vapor in equilibrium with its liquid.

• Factors: Intermolecular forces, temperature.

Solids

• Classified by bonding: ionic, molecular, covalent network, metallic.

• Properties: melting point, hardness, conductivity.

Gas Laws

• Ideal Gas Law: $PV = nRT$

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

• Gas stoichiometry: Relate volumes, moles, and masses in reactions involving gases.

Unit 8: Chemical Kinetics

Rate Laws and Rate Constants

• Rate law: $\text{rate} = k[A]^m[B]^n$

• Determine order from data; units of $k$ depend on overall order.

Factors Affecting Reaction Rate

• Temperature, concentration, surface area, and catalysts affect rate (collision theory).

Reaction Mechanisms

• Mechanism: Sequence of steps; slow step determines rate law.

• Identify intermediates and catalysts.

Unit 9: Chemical Equilibrium and Solubility

Equilibrium Expressions

• $K_{eq} = \frac{[\text{products}]}{[\text{reactants}]}$ (coefficients become exponents)

• ICE tables help solve for equilibrium concentrations.

Le Châtelier's Principle

• Predicts system response to changes in concentration, pressure, or temperature.

Solubility Product (Ksp)

• Relates to the solubility of sparingly soluble salts.

• Predict precipitation using $Q$ vs. $K_{sp}$.

Unit 10: Acids, Bases, and Titrations

Properties of Acids and Bases

• Acids: Donate H+; Bases: Accept H+ (Brønsted-Lowry).

• Strong vs. weak acids/bases: Degree of ionization.

pH and Ka Calculations

• $pH = -\log[H^+]$

• $K_a$ relates to acid strength; use ICE tables for weak acids.

• % dissociation = $\frac{[\text{dissociated}]}{[\text{initial}]} \times 100\%$

Titrations and Indicators

• Use titration curves to determine equivalence point.

• Select indicators based on expected pH at equivalence.

Unit 11: Electrochemistry

Electrochemical Cells

• Galvanic (voltaic) cells: Spontaneous redox reactions produce electricity.

• Label anode (oxidation), cathode (reduction), electron flow.

Cell Potentials and Faraday's Law

• $E_{cell} = E_{cathode} - E_{anode}$

• Faraday's Law: Relates charge, current, and amount of substance produced.

• $q = nF$ (where $n$ = moles of electrons, $F$ = Faraday's constant, $96,485$ C/mol e-)

Electrolysis

• Predict products of electrolysis for aqueous salt solutions with inert electrodes.

Additional info: This guide summarizes the core learning objectives and foundational concepts for a two-semester General Chemistry sequence, organized by unit for efficient exam preparation.