General Chemistry I: Comprehensive Study Guide (Chapters 1–13)

Chapter 1: Introduction to Chemistry – Matter, Energy, and Measurement

What is Matter? What is Chemistry?

• Matter is anything that has mass and occupies space.

• Chemistry is the study of the composition, structure, properties, and changes of matter.

Macroscopic vs. Submicroscopic (Molecular View)

• Macroscopic: Observable with the naked eye (e.g., ice, water, steam).

• Submicroscopic: At the atomic or molecular level, not directly observable (e.g., arrangement of H2O molecules).

Atoms vs. Molecules

• Atom: The smallest unit of an element that retains its chemical identity.

• Molecule: Two or more atoms chemically bonded together.

• Example: O2 is a molecule of two oxygen atoms; H2O is a molecule of two hydrogen and one oxygen atom.

States of Matter

• Solid, liquid, and gas differ in arrangement and motion of particles.

• Solids: Definite shape and volume; particles vibrate in place.

• Liquids: Definite volume, no definite shape; particles move past each other.

• Gases: No definite shape or volume; particles move freely.

Classification of Matter

• Pure Substances: Elements (cannot be broken down) and compounds (composed of two or more elements in fixed proportions).

• Mixtures: Homogeneous (uniform composition) and heterogeneous (non-uniform composition).

Physical vs. Chemical Properties and Changes

• Physical Properties: Observed without changing composition (e.g., melting point, density).

• Chemical Properties: Observed during a chemical change (e.g., flammability, reactivity).

• Physical Change: Does not alter composition (e.g., melting ice).

• Chemical Change: Alters composition (e.g., rusting iron).

Separation of Mixtures

• Physical methods: Filtration, distillation, chromatography.

• Chemical methods: Used to separate compounds into elements.

Measurement and Units

• SI Units: Mass (kg), length (m), time (s), temperature (K), amount (mol).

• Accuracy: Closeness to true value; Precision: Reproducibility of measurements.

• Significant Figures: Reflect precision in measurements.

Dimensional Analysis (Factor-Label Method)

• Used to convert between units using conversion factors.

• Example: To convert 12.9 kJ to J: $12.9\ \text{kJ} \times \frac{1000\ \text{J}}{1\ \text{kJ}} = 12900\ \text{J}$

Chapter 2: Atoms, Molecules, and Ions

Dalton’s Atomic Theory

• All matter is composed of atoms.

• Atoms of a given element are identical; atoms of different elements differ.

• Atoms combine in simple whole-number ratios to form compounds.

• Atoms are rearranged in chemical reactions, not created or destroyed.

Law of Conservation of Mass and Law of Constant Composition

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

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

Structure of the Atom

• Composed of protons (positive, in nucleus), neutrons (neutral, in nucleus), and electrons (negative, outside nucleus).

• Atomic number (Z): Number of protons; Mass number (A): Protons + neutrons.

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

Symbols and Notation

• Isotope notation: $^{A}_{Z}\text{X}$, where X is the element symbol.

Periodic Table

• Elements arranged by increasing atomic number.

• Groups (columns) and periods (rows) indicate similar properties and trends.

Formulas and Naming Compounds

• Ionic compounds: Metal + nonmetal; use charges to determine formula.

• Covalent compounds: Nonmetals; use prefixes (mono-, di-, tri-, etc.).

• Binary acids: H + nonmetal (e.g., HCl); Oxyacids: H + polyatomic ion (e.g., H2SO4).

Chapter 3: Chemical Reactions and Stoichiometry

Balancing Chemical Equations

• Law of conservation of mass: Same number of each atom on both sides of the equation.

• Steps: Write formulas, balance atoms one at a time, check work.

Types of Chemical Reactions

• Combination, decomposition, single replacement, double replacement, combustion.

Stoichiometry

• Relates quantities of reactants and products using balanced equations.

• Mole concept: 1 mole = $6.022 \times 10^{23}$ particles.

• Molar mass: Mass of 1 mole of a substance (g/mol).

• Limiting reactant: Reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical yield: Maximum amount of product possible; Actual yield: Amount actually obtained; Percent yield: $\frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100\%$

Chapter 4: Reactions in Aqueous Solution

Types of Aqueous Reactions

• Precipitation reactions: Formation of an insoluble solid.

• Acid-base reactions: Transfer of H+ ions.

• Redox reactions: Transfer of electrons.

Solubility Rules

• Used to predict whether a precipitate will form in a reaction.

Writing Ionic Equations

• Total ionic equation: Shows all ions present.

• Net ionic equation: Shows only the ions and molecules directly involved in the reaction.

Concentration Units

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

Chapter 5: Thermochemistry

Energy, Work, and Heat

• Energy: Capacity to do work or produce heat.

• Kinetic energy: Energy of motion; Potential energy: Stored energy.

• First Law of Thermodynamics: Energy is conserved; $\Delta E = q + w$

• Enthalpy (H): $\Delta H = q_p$ (heat at constant pressure).

Calorimetry

• Used to measure heat changes in chemical reactions.

• q = mcΔT, where m = mass, c = specific heat, ΔT = temperature change.

Hess’s Law

• The enthalpy change for a reaction is the same, regardless of the number of steps.

Chapter 6: Electronic Structure of Atoms

Nature of Light

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

• $c = \lambda \nu$

• Energy of a photon: $E = h\nu$

Atomic Models and Quantum Numbers

• Bohr model: Electrons in quantized orbits.

• Quantum mechanical model: Electrons in orbitals defined by quantum numbers (n, l, ml, ms).

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

Electron Configurations

• Describes the arrangement of electrons in an atom.

• Example: Oxygen (O): 1s2 2s2 2p4

Chapter 7: Periodic Properties of the Elements

Periodic Trends

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

• Ionization energy: Decreases down a group, increases across a period.

• Electron affinity, metallic character, ionic size, and other trends.

Chapter 8: Basic Concepts of Chemical Bonding

Ionic and Covalent Bonding

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

• Covalent bonds: Sharing of electrons between nonmetals.

• Lewis structures: Visual representations of bonding and lone pairs.

Bond Polarity and Electronegativity

• Electronegativity: Ability of an atom to attract electrons in a bond.

• Polar covalent bonds: Unequal sharing of electrons.

Resonance Structures

• Some molecules can be represented by two or more valid Lewis structures.

Chapter 9: Molecular Geometry and Bonding Theories

VSEPR Theory

• Predicts shapes of molecules based on electron pair repulsion.

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

Hybridization

Electron Groups	Hybridization	Bond Angle	Shape (no lone pairs)
2	sp	180°	Linear
3	sp2	120°	Trigonal planar
4	sp3	109.5°	Tetrahedral

Chapter 10: Gases

Properties of Gases

• Gases have indefinite shape and volume, are compressible, and fill their containers.

Gas Laws

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

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

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

• Ideal Gas Law: $PV = nRT$

Dalton’s Law of Partial Pressures

• Total pressure is the sum of partial pressures of all gases present: $P_{\text{total}} = P_1 + P_2 + \ldots$

Kinetic Molecular Theory

• Explains gas behavior based on motion of particles.

Real Gases and van der Waals Equation

• Accounts for intermolecular forces and finite molecular volume.

Chapter 11: Liquids and Intermolecular Forces

Types of Intermolecular Forces

• London dispersion forces, dipole-dipole interactions, hydrogen bonding.

Properties of Liquids

• Surface tension, viscosity, vapor pressure, boiling and melting points.

Phase Changes

• Melting, freezing, vaporization, condensation, sublimation, deposition.

Chapter 12: Solids and Modern Materials

Types of Solids

• Crystalline (ordered structure) vs. amorphous (disordered).

• Molecular, ionic, covalent network, and metallic solids.

Properties of Solids

• Melting point, conductivity, hardness, and solubility.

Chapter 13: Properties of Solutions

Solution Terminology

• Solute: Substance dissolved; Solvent: Substance doing the dissolving.

• Concentration units: Molarity (M), molality (m), percent by mass, ppm, ppb.

Solubility and Factors Affecting It

• Temperature, pressure, nature of solute and solvent.

• Henry’s Law: $S_g = kP_g$ (solubility of gas is proportional to pressure).

Colligative Properties

• Depend on number of solute particles: boiling point elevation, freezing point depression, osmotic pressure.

• Boiling point elevation: $\Delta T_b = K_b m$

• Freezing point depression: $\Delta T_f = K_f m$

Colloids

• Mixtures with large particles that do not settle out like true solutions.

How Soaps Work

• Soaps have hydrophobic and hydrophilic ends, allowing them to emulsify oils in water.

Additional info: These notes are structured to cover the main topics and subtopics outlined in the provided question set, with expanded academic context and examples for clarity and exam preparation.