CHEM 1001 Final Exam Study Guide: Core Concepts and Practice Topics

Chapter 1: Chemistry Basics – Matter and Measurement

Classification of Matter

• Pure Substances: Matter with a fixed composition; can be elements (single type of atom) or compounds (two or more elements chemically combined).

• Mixtures: Physical combinations of two or more substances. Homogeneous mixtures have uniform composition (e.g., saltwater), while heterogeneous mixtures do not (e.g., salad).

Periodic Table Reading

• Elements are organized by periods (rows) and groups (columns).

• Metals are typically on the left and center, non-metals on the right, and noble gases in Group 18.

Chemical vs. Physical Properties and Changes

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

• Chemical properties: Describe reactivity (e.g., flammability, acidity).

• Physical changes: Do not alter chemical identity (e.g., melting ice).

• Chemical changes: Produce new substances (e.g., rusting iron).

Significant Figures and Unit Conversions

• Rules for counting significant figures are essential for reporting measurements accurately.

• Unit conversions require multiplying by conversion factors (e.g., 1 in = 2.54 cm).

• Scientific notation expresses numbers as a product of a coefficient and a power of ten (e.g., $3.0 \times 10^2$).

Density, Mass, and Volume

• Density formula: $d = \frac{m}{v}$

• Used to solve for mass, volume, or density when two are known.

Chapter 2: Atoms and Radioactivity

Subatomic Particles

• Protons: Positive charge, found in nucleus.

• Neutrons: Neutral, found in nucleus.

• Electrons: Negative charge, found in electron cloud.

Atomic Number, Mass Number, and Symbolic Notation

• Atomic number (Z): Number of protons; identifies the element.

• Mass number (A): Protons + neutrons.

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

Isotopes

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

Radioactivity and Nuclear Emissions

• Radioactivity: Spontaneous emission of particles or energy from unstable nuclei.

• Types of emissions:

  • Alpha (α): $^4_2\text{He}$

  • Beta (β): $^0_{-1}\text{e}$

  • Gamma (γ): High-energy photon

  • Positron: $^0_{+1}\text{e}$

  • Neutron: $^1_0\text{n}$

Balanced Nuclear Equations

• Mass and atomic numbers must be conserved on both sides of the equation.

Half-Life Calculations

• Half-life (t1/2): Time for half of a radioactive sample to decay.

• Equation: $\text{isotope remaining} = \left(\frac{1}{2}\right)^n \times \text{starting amount}$, where $n = \frac{\text{total time}}{\text{half-life}}$

Chapter 3: Compounds – How Elements Combine

Electron Arrangements and the Octet Rule

• Electrons fill shells according to $2n^2$ rule (n = shell number).

• Octet rule: Atoms tend to have 8 electrons in their valence shell for stability.

Periods, Groups, and Electronic Configuration

• Period number: Highest energy level occupied.

• Group number: Number of valence electrons (old group numbers).

Formation of Ions

• Atoms become stable by gaining or losing electrons to become isoelectronic with the nearest noble gas.

• Cations: Lose electrons, positive charge (usually metals).

• Anions: Gain electrons, negative charge (usually non-metals).

Predicting Ionic Charges

• Main group elements form predictable charges based on group number (e.g., Group 1: +1, Group 17: –1).

Molar Mass and Avogadro’s Number

• Molar mass: Mass of one mole of a substance (g/mol), calculated from the periodic table.

• Avogadro’s number: $6.02 \times 10^{23}$ particles per mole.

Ionic and Covalent Bonds

• Ionic bonds: Transfer of electrons from metal to non-metal.

• Covalent bonds: Sharing of electrons between non-metals.

• Ionic compounds: Composed of ions, high melting points.

• Covalent compounds: Composed of molecules, lower melting points.

Chapter 4: Introduction to Organic Compounds

Structural Representations

• Lewis structures: Show all atoms and bonds.

• Condensed structures: Grouped atoms (e.g., CH3CH2CH3).

• Skeletal structures: Lines represent carbon chains; hydrogens often omitted.

Naming Alkanes and Cycloalkanes

• First ten alkanes: methane, ethane, propane, butane, pentane, hexane, heptane, octane, nonane, decane.

• Cycloalkanes: Ring structures; prefix "cyclo-" (e.g., cyclopentane).

General Formulas

• Alkane: $\text{C}_n\text{H}_{2n+2}$

• Cycloalkane: $\text{C}_n\text{H}_{2n}$

Functional Groups

• Specific groups of atoms that determine chemical properties (e.g., alcohols, ketones, carboxylic acids).

Chapter 5: Chemical Reactions

Thermodynamics and Reaction Types

• Exothermic: Releases heat.

• Endothermic: Absorbs heat.

• Enthalpy (H): Heat content.

• Entropy (S): Disorder or randomness.

• Free energy (G): Determines spontaneity ($\Delta G = \Delta H - T\Delta S$).

• Activation energy: Minimum energy required for reaction.

• Exergonic: Releases free energy.

• Endergonic: Absorbs free energy.

• Spontaneous: Occurs without external input.

• Non-spontaneous: Requires energy input.

Reaction Energy Diagrams

• Show energy changes during a reaction; labels include reactants, products, activation energy, and overall energy change.

Factors Affecting Reaction Rate

• Temperature (higher = faster), reactant concentration, and catalysts (lower activation energy).

Types of Chemical Reactions

• Synthesis: Two or more substances combine.

• Decomposition: One substance breaks down.

• Exchange/Displacement: Atoms or ions swap places.

Chapter 7: States of Matter and Their Attractive Forces

Properties of Gases

• Compressible, expand to fill container, low density.

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)

• Ideal Gas Law: $PV = nRT$

Attractive Forces

• Intermolecular forces: London dispersion, dipole-dipole, hydrogen bonding.

Oils and Fats

• Nonpolar, hydrophobic, differ in melting points and structure.

Chapter 8: Solution Chemistry

Solute and Solvent

• Solute: Substance dissolved.

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

Types of Mixtures

• Solution: Homogeneous, small particles (e.g., saltwater).

• Colloid: Medium particles, scatter light (e.g., milk).

• Suspension: Large particles, settle out (e.g., muddy water).

Solubility Factors

• Temperature (usually increases solubility of solids, decreases for gases).

• Pressure (mainly affects gases).

Electrolytes

• Strong electrolytes: Completely dissociate, good conductors (e.g., NaCl).

• Weak electrolytes: Partially dissociate, poor conductors (e.g., acetic acid).

• Non-electrolytes: Do not conduct (e.g., sugar).

Concentration Units and Dilution

• Molarity (M): $\text{mol/L}$

• Percent by mass (m/m): $\frac{\text{mass solute}}{\text{mass solution}} \times 100\%$

• Percent by volume (v/v): $\frac{\text{volume solute}}{\text{volume solution}} \times 100\%$

• Percent mass/volume (m/v): $\frac{\text{mass solute}}{\text{volume solution}} \times 100\%$

• Dilution equation: $C_{\text{initial}}V_{\text{initial}} = C_{\text{final}}V_{\text{final}}$

Chapter 9: Acids, Bases, and Buffers in the Body

Acids and Bases

• Acids: Donate H+ ions; sour taste, pH < 7.

• Bases: Accept H+ ions; bitter taste, pH > 7.

Strength of Acids and Bases

• Strong acids/bases: Completely ionize in water (e.g., HCl, NaOH).

• Weak acids/bases: Partially ionize (e.g., acetic acid, ammonia).

Neutralization Reactions

• Acid + base → salt + water.

Chemical Equilibrium

• Forward and reverse reactions occur at the same rate.

• Equilibrium constant expression: $K = \frac{[\text{products}]}{[\text{reactants}]}$ (exclude solids and liquids).

pH, Ka, and pKa

• pH: $-\log[\text{H}^+]$; measures acidity.

• Ka: Acid dissociation constant; higher Ka = stronger acid.

• pKa: $-\log K_a$; lower pKa = stronger acid.

pH Scale

• pH < 7: Acidic

• pH = 7: Neutral

• pH > 7: Basic