Comprehensive Review Notes for General, Organic, and Biological Chemistry Final Exam

Chapter 2: Chemistry and Measurements

Significant Figures and Scientific Notation

Accurate measurement and reporting are fundamental in chemistry. Significant figures reflect the precision of a measurement, while scientific notation is used to express very large or small numbers efficiently.

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

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

• Scientific Notation: Numbers are written as $a \times 10^n$, where $1 \leq a < 10$ and $n$ is an integer.

• Example: 0.00450 has three significant figures; 4.50 × 10-3 in scientific notation.

Units of Measurement and Dimensional Analysis

Chemistry uses the International System of Units (SI) for consistency. Dimensional analysis is a method to convert between units using conversion factors.

• Base SI Units: meter (m), kilogram (kg), second (s), mole (mol), kelvin (K), ampere (A), candela (cd).

• Dimensional Analysis: Multiply by conversion factors to cancel units.

• Example: To convert 25.0 cm to meters: $25.0\ \text{cm} \times \frac{1\ \text{m}}{100\ \text{cm}} = 0.250\ \text{m}$

Density and Specific Gravity

Density is the mass per unit volume of a substance. Specific gravity compares the density of a substance to that of water.

• Density Formula: $\text{Density} = \frac{\text{mass}}{\text{volume}}$

• Units: g/mL or g/cm3 for solids and liquids; g/L for gases.

• Specific Gravity: $\text{Specific Gravity} = \frac{\text{Density of substance}}{\text{Density of water}}$ (unitless)

• Example: A 10.0 g object with a volume of 2.00 mL has a density of $5.00\ \text{g/mL}$.

Chapter 3: Matter and Energy

Classification and Properties of Matter

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

• States of Matter: Solid (definite shape and volume), Liquid (definite volume, variable shape), Gas (variable shape and volume).

• Types of Matter:

  • Element: Pure substance of one type of atom (e.g., O2).

  • Compound: Substance of two or more elements chemically combined (e.g., H2O).

  • Mixture: Physical blend of two or more substances.

    • Homogeneous: Uniform composition (e.g., salt water).

    • Heterogeneous: Non-uniform composition (e.g., salad).

• Physical Change: Alters appearance, not composition (e.g., melting ice).

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

Energy and Heat

Energy is the capacity to do work. It is measured in joules (J) or calories (cal).

• Specific Heat: Amount of heat needed to raise the temperature of 1 g of a substance by 1°C.

  • Formula: $q = m \times c \times \Delta T$

  • Where $q$ = heat (J), $m$ = mass (g), $c$ = specific heat (J/g°C), $\Delta T$ = temperature change (°C)

• Heat of Fusion: Energy required to change 1 g of a solid to liquid at its melting point.

• Heating and Cooling Curves: Graphs showing temperature changes as heat is added or removed.

Chapter 4: Atoms and Elements

Atomic Structure and Subatomic Particles

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

• Dalton's Atomic Theory: Atoms are indivisible, each element has unique atoms, atoms combine in fixed ratios.

• Rutherford's Gold Foil Experiment: Showed atoms have a small, dense, positively charged nucleus.

• Subatomic Particles:

  • Proton: +1 charge, mass ≈ 1 amu

  • Neutron: 0 charge, mass ≈ 1 amu

  • Electron: -1 charge, mass ≈ 0.0005 amu

• Atomic Number (Z): Number of protons in the nucleus.

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

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

• Average Atomic Mass: Weighted average of isotopic masses.

• Electronic Arrangement: Electrons are arranged in energy levels (shells).

• Periodic Table: Elements arranged by increasing atomic number; groups (columns) and periods (rows).

• Metals, Nonmetals, Metalloids: Classified by properties (conductivity, luster, etc.).

• Valence Electrons: Electrons in the outermost shell, important for chemical bonding.

Chapter 5: Nuclear Chemistry

Types of Radiation and Nuclear Reactions

Nuclear chemistry studies changes in atomic nuclei, including radioactive decay and nuclear reactions.

• Types of Radiation: Alpha (α), Beta (β), Gamma (γ)

• Half-Life: Time required for half the nuclei in a sample to decay.

  • Formula: $N = N_0 \times (\frac{1}{2})^{t/t_{1/2}}$

• Nuclear Fission: Splitting a heavy nucleus into lighter nuclei, releasing energy.

• Nuclear Fusion: Combining light nuclei to form a heavier nucleus, releasing even more energy.

Chapter 6: Ionic and Molecular Compounds

Naming and Formulas of Compounds

Chemical compounds are named and written according to systematic rules.

• Ionic Compounds: Formed from metals and nonmetals; named by cation then anion (e.g., NaCl: sodium chloride).

• Polyatomic Ions: Charged groups of atoms (e.g., SO42-: sulfate).

• Covalent Compounds: Formed from nonmetals; use prefixes (e.g., CO2: carbon dioxide).

• Polar vs. Nonpolar Covalent: Polar bonds have unequal sharing of electrons; nonpolar have equal sharing.

• Molecular Shapes: Determined by VSEPR theory (e.g., linear, bent, tetrahedral).

Chapter 7: Chemical Quantities and Reactions

Chemical Equations and Stoichiometry

Chemical reactions are represented by balanced equations, which show the relationships between reactants and products.

• Writing and Balancing Equations: Law of conservation of mass requires equal numbers of each atom on both sides.

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

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

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

• Conversions: Moles ↔ grams using molar mass.

• Stoichiometry: Calculations based on balanced equations (mole-to-mole, mole-to-gram, etc.).

• Exothermic vs. Endothermic: Exothermic releases heat; endothermic absorbs heat.

Chapter 8: Gases

Kinetic Molecular Theory and Gas Laws

Gases are described by the kinetic molecular theory and obey several gas laws.

• Kinetic Molecular Theory: Gases consist of small particles in constant, random motion; collisions are elastic.

• Boyle's Law: $P_1V_1 = P_2V_2$ (at constant T and n)

• Molar Volume at STP: 1 mole of gas occupies 22.4 L at standard temperature and pressure (0°C, 1 atm).

• Sample Problem: If 2.00 mol of gas at STP, volume = $2.00 \times 22.4 = 44.8$ L.

Chapter 9: Solutions

Properties and Types of Solutions

Solutions are homogeneous mixtures of solute and solvent. Their properties depend on the nature of the components.

• Solute: Substance dissolved.

• Solvent: Substance doing the dissolving (often water).

• Types: Polar, nonpolar, strong/weak/non-electrolyte.

• Saturated/Unsaturated: Saturated holds maximum solute; unsaturated can dissolve more.

• Effect of Temperature: Solubility of solids increases with temperature; gases decrease.

• Concentration Units:

  • Mass/mass %: $\frac{\text{g solute}}{\text{g solution}} \times 100$

  • Mass/volume %: $\frac{\text{g solute}}{\text{mL solution}} \times 100$

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

  • Dilution: $M_1V_1 = M_2V_2$

• Osmosis and Tonicity: Isotonic (equal), hypertonic (higher solute), hypotonic (lower solute) solutions.

• Colloid, Suspension, Solution: Differ by particle size and stability.

Type	Particle Size	Appearance	Stability
Solution	<1 nm	Clear	Stable
Colloid	1-1000 nm	Cloudy	Stable
Suspension	>1000 nm	Cloudy, separates	Unstable

Chapter 10: Acids, Bases, and Equilibrium

Acids, Bases, and the Ionic Product of Water

Acids and bases are fundamental chemical species, and their behavior in water is described by equilibrium concepts.

• Acid: Donates H+ ions in water.

• Base: Accepts H+ ions or donates OH- ions.

• Ionic Product of Water: $K_w = [H^+][OH^-] = 1.0 \times 10^{-14}$ at 25°C

Chapter 11: Introduction to Organic Chemistry: Hydrocarbons

Types of Hydrocarbons

Hydrocarbons are organic compounds composed only of carbon and hydrogen.

• Alkanes: Saturated hydrocarbons (single bonds), general formula $C_nH_{2n+2}$.

• Alkenes: Unsaturated hydrocarbons (one or more double bonds), general formula $C_nH_{2n}$.

• Alkynes: Unsaturated hydrocarbons (one or more triple bonds), general formula $C_nH_{2n-2}$.

• Aromatic Compounds: Contain benzene ring structure.

• Isomers: Compounds with same molecular formula but different structures.

• Cycloalkanes: Ring-shaped alkanes, general formula $C_nH_{2n}$.

Chapter 13: Carbohydrates

Types and Naming of Carbohydrates

Carbohydrates are essential biomolecules classified by the number of sugar units.

• Monosaccharides: Single sugar units (e.g., glucose, fructose).

• Disaccharides: Two monosaccharides linked (e.g., sucrose, lactose).

• Polysaccharides: Many monosaccharides linked (e.g., starch, cellulose, glycogen).

Chapter 15: Lipids

Types of Lipids

Lipids are hydrophobic biomolecules with diverse structures and functions.

• Types: Fatty acids, triglycerides, phospholipids, steroids, waxes.

• Functions: Energy storage, cell membrane structure, signaling.

Chapter 16: Amino Acids, Proteins, and Enzymes

Classification and Structure of Proteins

Proteins are polymers of amino acids with complex structures and vital biological roles.

• Amino Acids: Building blocks of proteins; contain amino and carboxyl groups.

• Dipeptides: Two amino acids linked by a peptide bond.

• Protein Structure Levels:

  • Primary: Sequence of amino acids.

  • Secondary: Alpha-helix or beta-sheet folding.

  • Tertiary: 3D folding of a single polypeptide.

  • Quaternary: Association of multiple polypeptides.

• Enzymes: Biological catalysts that speed up reactions.

• Denaturation: Loss of protein structure and function due to heat, pH, or chemicals.