General Chemistry I: Comprehensive Study Guide

Physical vs. Chemical Changes

Definitions and Distinctions

• Physical changes: Changes in form or state without altering chemical composition (e.g., phase changes such as melting, boiling).

• Chemical changes: Changes that result in new substances being formed; involve breaking and forming chemical bonds (e.g., combustion, rusting).

• Key distinction: Physical changes are usually reversible; chemical changes are often irreversible.

• Example: Boiling water (physical), burning wood (chemical).

Intermolecular Forces vs. Intramolecular Bonds

Types of Chemical Interactions

• Intermolecular forces: Attractions between molecules, such as hydrogen bonding, dipole-dipole, and London dispersion forces. These affect physical properties like boiling and melting points.

• Intramolecular bonds: Bonds within molecules, including covalent, ionic, and metallic bonds. These are broken during chemical reactions.

• Example: Water molecules held together by hydrogen bonds (intermolecular); H–O bonds within a water molecule (intramolecular).

Gas Laws and Properties

Avogadro's Law

• Relates volume of a gas to the amount (moles) of gas at constant temperature and pressure.

• Avogadro's Law: $V \\propto n$ (at constant T and P)

• Example: Doubling the moles of gas doubles the volume.

Ideal Gas Law

• Relates pressure, volume, temperature, and moles of gas.

• Equation: $PV = nRT$

• Standard conditions: STP (0°C, 1 atm), molar volume = 22.4 L/mol.

Combined Gas Law

• Combines Boyle's, Charles's, and Gay-Lussac's laws.

• Equation: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

• Used when temperature, pressure, and volume all change.

Dalton's Law of Partial Pressures

• Total pressure of a mixture of gases equals the sum of the partial pressures.

• Equation: $P_{total} = P_1 + P_2 + P_3 + \ldots$

Kinetic Molecular Theory

• Describes behavior of gases in terms of particle motion.

• Average kinetic energy depends only on temperature: $KE_{avg} = \frac{3}{2}kT$

• Relationship between molecular mass and speed: $u_{rms} = \sqrt{\frac{3RT}{M}}$

Stoichiometry with Gases

• Use balanced equations to relate moles and volumes of gases at same conditions.

• Example: $2H_2 + O_2 \rightarrow 2H_2O$ (2 volumes H2 react with 1 volume O2).

Measurements and Significant Figures

Accuracy vs. Precision

• Accuracy: Closeness to true value.

• Precision: Closeness of repeated measurements to each other.

• Can be precise without being accurate.

Significant Figures

• Rules for counting significant digits in measurements.

• Rules for calculations:

  • Multiplication/division: Least number of significant figures.

  • Addition/subtraction: Least number of decimal places.

• Report answers with correct significant figures.

Unit Conversions

• Dimensional analysis: Use conversion factors to change units.

• Convert between metric units (e.g., cm to m).

• Area and volume conversions (e.g., cm2 to m2).

Density Calculations

• Formula: $d = \frac{m}{V}$

• Used to relate mass and volume of substances.

Atomic Structure and Composition

Key Experiments

• Millikan's oil drop experiment: Determined electron charge.

• Rutherford's gold foil experiment: Discovered nucleus and atomic structure.

Atomic Composition

• Atoms consist of protons, neutrons, and electrons.

• Atomic number (Z): Number of protons.

• Mass number (A): Protons + neutrons.

• Ion charges: Electron count differs from proton count.

Isotopes

• Atoms of the same element with different mass numbers.

• Calculate average atomic mass using isotope abundances.

• Notation: Element–mass number (e.g., C-12).

Nuclear Stability

• Factors: Neutron/proton ratio, nuclear forces.

• Stable nuclides vs. unstable (radioactive) nuclides.

The Mole and Stoichiometry

The Mole Concept

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

• Convert between mass, moles, and number of particles.

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

Percent Composition

• Calculate mass percent of each element in a compound.

• Formula: $\%\,\text{element} = \frac{\text{mass of element}}{\text{mass of compound}} \times 100$

Stoichiometry

• Use balanced equations to relate reactants and products.

• Mole ratios from coefficients.

• Limiting reactant problems: Identify which reactant limits product formation.

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

Light and Atomic Structure

Electromagnetic Radiation

• Wave properties: Wavelength ($\lambda$), frequency ($\nu$), energy ($E$).

• Relationship: $c = \lambda \nu$ and $E = h \nu$

• Inverse relationship between wavelength and energy.

Energy of Photons

• Calculate energy per photon: $E = h \nu$

• Calculate energy per mole: $E_{mol} = N_A h \nu$

Atomic Spectra

• Line spectra: Discrete energy levels.

• Energy absorbed when electron moves to higher level; released when moving to lower level.

• Larger energy gaps = larger energy changes.

Electron Transitions

• Calculate energy differences between levels.

• Identify transitions with most/least energy.

Quantum Numbers

• n (principal): Energy level (1, 2, ...)

• l (angular momentum): Subshell (0 to n-1)

• ml (magnetic): Orbital orientation (-l to +l)

• ms (spin): Electron spin (+1/2 or -1/2)

Orbitals

• Number of orbitals in each subshell: s(1), p(3), d(5), f(7)

• Maximum number of orbitals for given n: $n^2$

Electron Configuration

• Aufbau principle, Hund's rule, Pauli exclusion principle.

• Noble gas notation for shorthand.

• Exceptions: Cu, Ag, Cr.

• Electron configuration of ions: Remove electrons from highest n first.

Unpaired Electrons

• Determine from electron configuration.

• Important for magnetic properties.

Periodic Trends

Organization of Periodic Table

• Ordered by atomic number.

• Groups: Same number of valence electrons, similar properties.

• Periods: Same highest energy level.

Atomic Radius

• Increases down a group, decreases across a period.

Ionization Energy

• Energy required to remove an electron.

• Increases across a period, decreases down a group.

• Successive ionizations require more energy.

Electron Affinity

• Energy change when adding an electron.

• Most favorable for halogens, least for noble gases.

• Generally becomes more negative across a period.

Metallic Character

• Metals, nonmetals, metalloids: Group classifications and properties.

Chemical Bonding and Lewis Structures

Ionic Bonding

• Transfer of electrons from metal to nonmetal.

• Lattice energy: Higher charge and smaller size = higher lattice energy.

Covalent Bonding

• Sharing of electrons between nonmetals.

• Bond order: Single < double < triple.

• Bond strength: Single < double < triple.

• Bond length: Single > double > triple.

Resonance

• Multiple valid Lewis structures.

• Identify correct resonance structures.

• Equal contributors vs. major/minor forms.

Molecular Geometry (VSEPR)

Electron Domain Geometry (EDG)

• Based on all electron domains (bonding + lone pairs).

• Common geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Molecular Geometry (MG)

• Based only on positions of atoms.

• Effect of lone pairs on shape.

Bond Angles

• Ideal angles for each geometry.

• Effect of lone pairs: Compression of bond angles.

Predicting Geometries

• Count electron domains around central atom.

• Apply VSEPR theory.

Polarity and Intermolecular Forces

Molecular Polarity

• Depends on bond polarity and molecular geometry.

• Symmetrical molecules with polar bonds can be nonpolar overall.

Intermolecular Forces

• Relative strength: London dispersion < dipole-dipole < hydrogen bonding.

• Stronger IMF = higher boiling point.

• Compare boiling points based on IMF types and molecular size.

Hybridization and Molecular Orbital Theory

Hybridization

• Determine from electron domain geometry.

• Types: sp (linear), sp2 (trigonal planar), sp3 (tetrahedral), sp3d (trigonal bipyramidal), sp3d2 (octahedral).

Sigma and Pi Bonds

• σ bonds: End-to-end overlap; all single bonds.

• π bonds: Side-to-side overlap; in double and triple bonds.

• Single bond = 1σ, double bond = 1σ + 1π, triple bond = 1σ + 2π.

Nomenclature and Formulas

Ionic Compounds

• Metal + nonmetal.

• Transition metals: Use Roman numerals for charge.

• Polyatomic ions: Memorize common ions.

Molecular Compounds

• Nonmetal + nonmetal.

• Prefixes: mono-, di-, tri-, etc.

Acids

• -ate becomes -ic acid; -ite becomes -ous acid.

• Prefixes: per-, hypo-.

Common Polyatomic Ions

Ion	Formula
Nitrate	NO3-
Sulfate	SO42-
Carbonate	CO32-
Phosphate	PO43-
Hydroxide	OH-
Cyanide	CN-
Acetate	CH3COO-

Chemical Reactions and Equations

Balancing Equations

• Conservation of mass: Atoms must balance on both sides.

• Balance atoms first, then coefficients.

Reaction Types

• Combination/Synthesis

• Decomposition

• Single Replacement

• Double Replacement

• Combustion

Predicting Products

• Solubility rules for precipitation reactions.

• Acid-base neutralization.

• Gas-forming reactions.

Solutions and Solution Stoichiometry

Concentration Units

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

• Calculate molarity from mass and volume.

Dilution

• Equation: $M_1V_1 = M_2V_2$

• Increasing volume decreases concentration.

Solution Stoichiometry

• Use molarity to find moles.

• Apply stoichiometric ratios.

• Calculate volumes needed for reactions.

Acids, Bases, and pH

Acids and Bases

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

• Strong vs. weak acids and bases.

• Substances that act as bases in water.

pH Scale

• Equation: $pH = -\log[H^+]$

• pH < 7: acidic; pH = 7: neutral; pH > 7: basic.

• Calculate pH from [H+] and vice versa.

Titrations

• Neutralization reactions.

• Calculate volumes needed using stoichiometry.

Oxidation-Reduction (Redox)

Oxidation States

• Rules for assigning oxidation numbers.

• Change in oxidation state indicates redox reaction.

Oxidation and Reduction

• Oxidation: Loss of electrons (increase in oxidation state).

• Reduction: Gain of electrons (decrease in oxidation state).

• OIL RIG: Oxidation Is Loss, Reduction Is Gain.

Oxidizing and Reducing Agents

• Oxidizing agent: Causes oxidation; itself reduced.

• Reducing agent: Causes reduction; itself oxidized.

Formal Charge vs. Oxidation State

• Formal charge: Used for Lewis structures.

• Oxidation state: Used for redox reactions.

Activity Series

• More active metals are more easily oxidized.

• More active metal can displace less active metal from solution.

Thermochemistry

Energy Units

• Temperature changes and phase changes.

Enthalpy ($\Delta H$)

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

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

Calorimetry

• Coffee-cup calorimeter: Constant pressure, measures $q_{rxn}$.

• Bomb calorimeter: Constant volume, $q = C \Delta T$ (C is heat capacity).

• Calculate $\Delta H$ per mole from experimental data.

Hess's Law

• Enthalpy is a state function.

• $\Delta H_{rxn} = \Sigma \Delta H_f^{\circ}$(products) $- \Sigma \Delta H_f^{\circ}$(reactants)

• Manipulate equations and enthalpies to find desired $\Delta H$.

Standard Enthalpy of Formation ($\Delta H_f^{\circ}$)

• Enthalpy change to form 1 mole from elements in standard states.

• Elements in standard state: $\Delta H_f^{\circ} = 0$

• Calculate $\Delta H_{rxn}$ using $\Delta H_f^{\circ}$ values.

Stoichiometry with Energy

• Use molar enthalpy values with stoichiometric ratios.

• Calculate energy released/absorbed for given amounts.

Key Formulas to Know

• Density: $d = \frac{m}{V}$

• Heat: $q = mc\Delta T$

• Calorimetry: $q = C\Delta T$

• Hess's Law: $\Delta H_{rxn} = \Sigma \Delta H_f^{\circ}(\text{products}) - \Sigma \Delta H_f^{\circ}(\text{reactants})$

• Graham's Law: $\text{Rate}_1/\text{Rate}_2 = \sqrt{M_2/M_1}$

• Partial Pressure: $P_A = X_A \times P_{total}$

Study Tips

1. Practice drawing Lewis structures and predicting geometries.

2. Memorize polyatomic ions and solubility rules.

3. Work through stoichiometry problems step-by-step.

4. Understand trends rather than memorizing specific values.

5. Practice unit conversions and dimensional analysis.

6. Review electron configurations and orbital diagrams.

7. Understand the relationships between gas law variables.

8. Practice pH and molarity calculations.

9. Work through enthalpy problems using Hess's Law.

10. Review oxidation states and redox identification.