CHEM 1025 Final Exam Review Topics (Ch 2-10 & 13): Structured Study Notes

• Chapter 2: Measurement and Problem Solving

Scientific Measurement

Measurement is fundamental in chemistry for quantifying substances and reactions. It involves using units and tools to obtain reliable data.

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit. They reflect the precision of a measurement.

• Basic Units of Measurement: The SI system includes meters (length), kilograms (mass), seconds (time), and moles (amount of substance).

• Unit Conversions: Changing from one unit to another using conversion factors.

• Density: The ratio of mass to volume, calculated as $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$.

Example: Converting 10.0 cm to meters: $10.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.100\ \text{m}$

Chapter 3: Matter and Energy

Matter Classification

Matter is anything that has mass and occupies space. It can be classified by composition and properties.

• Pure Substances: Elements and compounds with fixed composition.

• Mixtures: Physical combinations of substances. Homogeneous mixtures are uniform throughout; heterogeneous mixtures are not.

• Physical vs. Chemical Changes: Physical changes alter appearance, not composition; chemical changes produce new substances.

• Energy: The capacity to do work. Includes kinetic (motion) and potential (stored) energy.

Example: Dissolving salt in water is a physical change; burning wood is a chemical change.

Chapter 4: Atoms and Elements

Atomic Theory and Structure

Atoms are the basic units of matter, composed of protons, neutrons, and electrons.

• Dalton's Atomic Theory: Matter is made of atoms, which combine in fixed ratios to form compounds.

• Thomson and Rutherford Models: Thomson discovered electrons; Rutherford described the nucleus.

• Periodic Table: Organizes elements by atomic number and properties. Groups (columns) share chemical behavior; periods (rows) indicate energy levels.

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

Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Chapter 5: Molecules and Compounds

Chemical Bonding and Formulas

Compounds are formed when atoms bond together. Chemical formulas show the types and numbers of atoms.

• Ionic Compounds: Formed from metals and nonmetals; consist of ions held by electrostatic forces.

• Molecular Compounds: Formed from nonmetals; consist of molecules with covalent bonds.

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

Example: Water's molecular formula is H2O; its empirical formula is also H2O.

Chapter 6: Chemical Composition

Quantifying Compounds

Chemists use the mole to count particles and relate mass to number of atoms or molecules.

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

• Molar Mass: Mass of one mole of a substance, in grams per mole.

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

Example: Calculate moles in 18 g of H2O: $\text{Moles} = \frac{18\ \text{g}}{18.02\ \text{g/mol}} = 1.00\ \text{mol}$

Chapter 7: Chemical Reactions

Types and Representation of Reactions

Chemical reactions involve the transformation of substances via breaking and forming bonds.

• Balancing Equations: Ensures the same number of atoms of each element on both sides.

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

• Precipitation and Acid-Base Reactions: Formation of insoluble products or transfer of protons.

Example: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

Chapter 8: Quantities in Chemical Reactions

Stoichiometry

Stoichiometry uses balanced equations to calculate amounts of reactants and products.

• Limiting Reactant: The reactant that is completely consumed first, limiting product formation.

• Theoretical Yield: Maximum amount of product possible.

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

Example: If 10 g product is obtained but 12 g is possible, percent yield is $\frac{10}{12} \times 100\% = 83.3\%$

Chapter 9: Electrons in Atoms and the Periodic Table

Atomic Structure and Electron Configuration

Electron arrangement determines chemical properties and periodic trends.

• Orbitals: Regions where electrons are likely found; s, p, d, f types.

• Electron Configuration: Distribution of electrons among orbitals, e.g., $1s^2 2s^2 2p^6$.

• Periodic Trends: Atomic radius, ionization energy, and electronegativity vary across the table.

Example: Sodium: $1s^2 2s^2 2p^6 3s^1$

Chapter 10: Chemical Bonding

Types of Chemical Bonds

Atoms bond to achieve stable electron configurations, forming molecules and compounds.

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

• Covalent Bonds: Sharing of electrons between nonmetals.

• Lewis Structures: Diagrams showing bonding and lone pairs.

• Resonance: Some molecules have multiple valid Lewis structures.

Example: CO2 has two resonance structures.

Chapter 13: Solutions

Properties and Calculations for Solutions

Solutions are homogeneous mixtures of solute and solvent. Their properties depend on concentration and interactions.

• Solubility: Maximum amount of solute that dissolves in solvent at a given temperature.

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

• Types of Solutions: Saturated, unsaturated, and supersaturated.

• Preparing Solutions: Calculating mass or volume needed for desired concentration.

Example: To prepare 1 L of 0.5 M NaCl, dissolve 29.2 g NaCl in water.

Additional Info

• Periodic Table: Reference for atomic numbers, symbols, and masses.

• Review all definitions, equations, and examples for exam preparation.