General Chemistry Final Exam Review: Key Concepts and Calculations

Chapter 2: Measurements and Units

Metric Conversions

Understanding metric conversions is essential for accurate scientific measurements. The metric system uses prefixes to indicate powers of ten.

• Key Prefixes: kilo- (103), centi- (10-2), milli- (10-3), micro- (10-6).

• Conversion Example: To convert 5.0 cm to meters: $5.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.050\ \text{m}$

Types of Measurements and Units

• Base SI Units: meter (length), kilogram (mass), second (time), mole (amount of substance).

• Derived Units: volume (liter), density (g/cm3).

Significant Figures and Rounding

• Significant Figures: Digits that carry meaning in a measurement.

• Rules: Nonzero digits are always significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant only if there is a decimal point.

• Example: 0.00450 has three significant figures.

Dimensional Analysis

• Definition: A method to convert units using conversion factors.

• Formula: $\text{Quantity in new units} = \text{Original quantity} \times \frac{\text{Desired unit}}{\text{Original unit}}$

Density

• Definition: Mass per unit volume.

• Formula: $\text{Density} = \frac{\text{Mass}}{\text{Volume}}$

• Example: If a block has a mass of 10 g and a volume of 2 cm3, its density is $5\ \text{g/cm}^3$.

Chapter 3: Matter and Its Properties

Types of Matter

• Pure Substances: Elements and compounds.

• Mixtures: Homogeneous (solutions) and heterogeneous (suspensions).

Chemical and Physical Properties

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

• Chemical Properties: Describe ability to undergo chemical change (e.g., flammability).

Specific Heat

• Definition: Amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Formula: $q = m c \Delta T$

• Example: Calculate heat needed to raise 10 g of water by 5°C ($c = 4.18\ \text{J/g}^\circ\text{C}$): $q = 10 \times 4.18 \times 5 = 209\ \text{J}$

Chapter 4: Atoms and Elements

Element Names and Symbols

• Periodic Table: Each element has a unique symbol (e.g., H for hydrogen, O for oxygen).

Isotopes

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

• Example: Carbon-12 and Carbon-14.

Atomic Structure Calculations

• Atomic Number (Z): Number of protons.

• Mass Number (A): Number of protons + neutrons.

• Example: For $^{14}_{6}\text{C}$, Z = 6, A = 14.

Chapter 5: Electrons and Periodic Trends

Electromagnetic Spectrum

• Definition: Range of all types of electromagnetic radiation.

• Order: Radio, microwave, infrared, visible, ultraviolet, X-ray, gamma ray.

Energy Levels, Sublevels, Orbitals

• Energy Levels: Principal quantum number (n).

• Sublevels: s, p, d, f.

• Orbitals: Regions of space where electrons are likely found.

Electron Configurations

• Aufbau Principle: Electrons fill lowest energy orbitals first.

• Example: Oxygen: 1s2 2s2 2p4

Periodic Trends

• Trends: Atomic radius decreases across a period, increases down a group; ionization energy increases across a period.

Chapter 6: Ions and Compounds

Cations, Anions, Ionic Charge

• Cation: Positively charged ion (loss of electrons).

• Anion: Negatively charged ion (gain of electrons).

Naming Ionic and Molecular Compounds

• Ionic Compounds: Name cation first, then anion (e.g., NaCl: sodium chloride).

• Molecular Compounds: Use prefixes (e.g., CO2: carbon dioxide).

Chapter 7: Chemical Quantities

Avogadro's Number and Molar Mass

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

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

Mole-to-Mole Ratios

• Definition: Ratio of moles of reactants and products in a balanced equation.

Mass Percent

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

Chapter 8: Chemical Reactions

Reaction Types

• Synthesis, Decomposition, Single Replacement, Double Replacement, Combustion.

Balancing Reactions

• Law of Conservation of Mass: Number of atoms of each element must be equal on both sides.

Oxidation-Reduction (Redox)

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Identify: Assign oxidation numbers to determine which species is oxidized/reduced.

Chapter 9: Stoichiometry and Limiting Reactants

Stoichiometric Calculations

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

• Example: $2H_2 + O_2 \rightarrow 2H_2O$

Limiting Reactant

• Definition: Reactant that is completely consumed first, limiting the amount of product.

Percent Yield

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

Chapter 10: Molecular Structure and Interactions

Lewis Structures

• Drawing: Show valence electrons as dots; connect atoms to satisfy octet rule.

Molecular Shape and Bond Polarity

• VSEPR Theory: Predicts 3D shape based on electron pair repulsion.

• Bond Polarity: Difference in electronegativity between atoms.

Intermolecular Interactions

• Types: Hydrogen bonding, dipole-dipole, London dispersion forces.

Heating/Cooling Curves

• Shows: Temperature change as heat is added/removed; plateaus indicate phase changes.

Chapter 11: Gases and Gas Laws

Kinetic Molecular Theory

• Describes: Behavior of gases in terms of particle motion.

Factors Affecting Gases

• Pressure, Volume, Temperature, Number of Moles.

Units of Pressure

• Common Units: atm, mmHg, torr, Pa.

Gas Law Formulas

• Ideal Gas Law: $PV = nRT$

• STP: Standard Temperature (0°C) and Pressure (1 atm).

Stoichiometric Calculations with Gases

• Use Ideal Gas Law to relate moles, volume, and pressure in reactions.

Chapter 12: Solutions

Definitions and Vocabulary

• Solution: Homogeneous mixture of solute and solvent.

• Concentration: Amount of solute per unit volume of solution.

Calculations of Concentration

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

Dilutions

• Formula: $M_1V_1 = M_2V_2$

Chapter 13: Chemical Kinetics and Equilibrium

Reaction Requirements

• Reactants must collide with sufficient energy and proper orientation.

Factors Affecting Rate

• Concentration, temperature, catalysts, surface area.

Equilibrium

• Definition: State where forward and reverse reaction rates are equal.

• Equilibrium Constant (K): $K = \frac{[\text{products}]}{[\text{reactants}]}$

Le Châtelier's Principle

• System shifts to counteract changes in concentration, pressure, or temperature.

Chapter 14: Acids and Bases

Identifying Acids and Bases

• Acids: Donate H+ ions.

• Bases: Accept H+ ions or donate OH-.

Strong and Weak Acids/Bases

• Strong: Completely dissociate in water.

• Weak: Partially dissociate.

Conjugate Acid-Base Pairs

• Definition: Acid and base that differ by one proton.

pH Calculations

• Formula: $\text{pH} = -\log[H^+]$

• Relationship: $[H^+][OH^-] = 1 \times 10^{-14}$

Buffer Systems

• Definition: Solutions that resist changes in pH.

Titration Calculations

• Used to determine concentration of an acid or base using a neutralization reaction.