General Chemistry I: Comprehensive Study Guide (Chapters 1–11)

Chapter 1: Matter, Measurement & Problem Solving

Essential Concepts and Symbols

This chapter introduces the foundational concepts of chemistry, including the nature of matter, measurement techniques, and the scientific method.

• Matter: Anything that has mass and occupies space. Classified as elements, compounds, or mixtures.

• Atoms and Molecules: Atoms are the basic units of matter; molecules are combinations of atoms bonded together.

• Chemical Symbols: Shorthand representations of elements (e.g., H for hydrogen).

• Scientific Method: Systematic approach to knowledge involving observation, hypothesis, experimentation, and theory development.

• Physical vs. Chemical Changes: Physical changes alter appearance without changing composition; chemical changes result in new substances.

• Properties of Matter: Includes physical (e.g., melting point) and chemical (e.g., reactivity) properties.

• Units of Measurement: SI units (meter, kilogram, second, mole, etc.) are standard in chemistry.

• Accuracy and Precision: Accuracy refers to closeness to true value; precision refers to reproducibility.

• Dimensional Analysis: Technique for converting between units using conversion factors.

Example: Converting 25.0 cm to meters using dimensional analysis: $25.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.250\ \text{m}$

Chapter 2: Atoms & Elements

Atomic Structure and the Nature of Elements

This chapter explores the structure of atoms, the development of atomic theory, and the organization of elements.

• Dalton's Atomic Theory: Matter is composed of atoms; atoms of each element are identical; atoms combine in simple ratios to form compounds.

• Subatomic Particles: Protons (+), neutrons (0), electrons (-).

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

• Atomic Mass: Weighted average mass of an element's isotopes.

• Periodic Table: Arrangement of elements by increasing atomic number; groups/families share chemical properties.

• Law of Definite Proportions: Compounds contain elements in fixed ratios.

• Law of Multiple Proportions: Elements can combine in different ratios to form different compounds.

Example: Carbon-12 and Carbon-13 are isotopes of carbon with atomic masses of 12 and 13 amu, respectively.

Chapter 3: Molecules and Compounds

Chemical Bonding and Compound Formation

This chapter covers the formation and properties of molecules and compounds, including ionic and covalent bonding.

• Chemical Bonds: Ionic (transfer of electrons) and covalent (sharing of electrons).

• Formulas: Empirical (simplest ratio), molecular (actual number of atoms), and structural (arrangement of atoms).

• Naming Compounds: Systematic rules for naming ionic and molecular compounds.

• Mole Concept: 1 mole = $6.022 \times 10^{23}$ entities (Avogadro's number).

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

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

Example: Water ($H_2O$) has a molar mass of $18.02\ \text{g/mol}$.

Chapter 4: Chemical Reactions and Chemical Quantities

Balancing and Classifying Chemical Equations

This chapter focuses on chemical reactions, balancing equations, and stoichiometric calculations.

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

• Balancing Equations: Ensuring equal numbers of each atom on both sides of the equation.

• Stoichiometry: Quantitative relationships between reactants and products.

• Limiting Reactant: The reactant that determines the amount of product formed.

• Theoretical Yield: Maximum amount of product possible.

• Percent Yield: $\frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$

Example: In the reaction $2H_2 + O_2 \rightarrow 2H_2O$, 2 moles of $H_2$ react with 1 mole of $O_2$ to produce 2 moles of water.

Chapter 5: Introduction to Solutions and Aqueous Solutions

Solution Concentration and Stoichiometry

This chapter introduces solutions, their concentrations, and reactions in aqueous media.

• Concentration Units: Molarity ($M = \frac{\text{moles solute}}{\text{liters solution}}$), molality, percent by mass.

• Solution Stoichiometry: Calculations involving concentrations and volumes.

• Types of Aqueous Solutions: Electrolytes (conduct electricity) and nonelectrolytes.

• Precipitation Reactions: Formation of insoluble products.

• Acid-Base and Redox Reactions: Transfer of protons or electrons in solution.

Example: Mixing $NaCl$ and $AgNO_3$ forms a precipitate of $AgCl$.

Chapter 6: Gases

Properties and Laws of Gases

This chapter examines the behavior of gases and the laws governing their properties.

• Kinetic Molecular Theory: Explains gas behavior based on particle motion.

• Gas Laws:

  • Boyle's Law: $P_1V_1 = P_2V_2$ (pressure and volume are inversely related)

  • Charles's Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (volume and temperature are directly related)

  • Ideal Gas Law: $PV = nRT$

• Partial Pressure: Pressure exerted by each gas in a mixture.

• Deviations from Ideal Behavior: Real gases deviate at high pressure and low temperature.

Example: Calculate the volume occupied by 1 mole of an ideal gas at STP: $V = \frac{nRT}{P}$

Chapter 7: Thermochemistry

Energy Changes in Chemical Reactions

This chapter explores the concepts of energy, heat, and work in chemical processes.

• Types of Energy: Kinetic (motion), potential (position), thermal (heat).

• First Law of Thermodynamics: Energy cannot be created or destroyed; $\Delta E = q + w$

• Enthalpy ($\Delta H$): Heat change at constant pressure.

• Calorimetry: Measurement of heat changes using calorimeters.

• Specific Heat Capacity ($C$): Amount of heat required to raise temperature of 1 g by 1°C.

• Hess's Law: Total enthalpy change is the sum of enthalpy changes for individual steps.

Example: $q = mC\Delta T$ calculates heat absorbed or released.

Chapter 8: The Quantum-Mechanical Model of the Atom

Atomic Structure and Quantum Theory

This chapter introduces quantum mechanics and its application to atomic structure.

• Electromagnetic Radiation: Light as oscillating electric and magnetic fields.

• Key Equations:

  • $c = \lambda \nu$ (speed of light = wavelength × frequency)

  • $E = h\nu$ (energy of a photon = Planck's constant × frequency)

• Bohr Model: Electrons occupy quantized energy levels.

• Quantum Numbers: Describe electron energy, shape, orientation, and spin.

• Wave-Particle Duality: Electrons exhibit both wave and particle properties.

Example: Calculate energy of a photon with frequency $\nu$: $E = h\nu$

Chapter 9: Periodic Properties of the Elements

Trends in the Periodic Table

This chapter discusses periodic trends and their explanations.

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

• Ionization Energy: Energy required to remove an electron; increases across a period.

• Electron Affinity: Energy change when an atom gains an electron.

• Electronegativity: Tendency to attract electrons in a bond.

Example: Fluorine has the highest electronegativity in the periodic table.

Chapter 10: Chemical Bonding I: The Lewis Model

Lewis Structures and Bonding

This chapter covers the representation of molecules using Lewis structures and the basics of chemical bonding.

• Ionic vs. Covalent Bonds: Ionic bonds involve electron transfer; covalent bonds involve sharing.

• Lewis Structures: Diagrams showing valence electrons and bonds.

• Octet Rule: Atoms tend to have eight electrons in their valence shell.

• Exceptions to Octet Rule: Some elements have fewer or more than eight electrons.

• Bond Polarity: Difference in electronegativity leads to polar bonds.

• Formal Charge: $\text{Formal Charge} = \text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$

Example: Draw the Lewis structure for $CO_2$.

Chapter 11: Chemical Bonding II: Molecular Shapes, VSEPR & MO Theory

Molecular Geometry and Bonding Theories

This chapter explains how molecular shapes are determined and introduces advanced bonding theories.

• VSEPR Theory: Valence Shell Electron Pair Repulsion theory predicts molecular shapes.

• Molecular Polarity: Determined by shape and bond polarity.

• Hybridization: Mixing of atomic orbitals to form new hybrid orbitals.

• Molecular Orbital Theory: Electrons are delocalized over the entire molecule.

Example: $CH_4$ has a tetrahedral geometry due to $sp^3$ hybridization.

Additional info: These notes are based on a study guide outline for a General Chemistry I final exam, covering chapters 1–11. The content has been expanded and organized for clarity and completeness.