CHM101 Final Exam Review: Key Concepts and Study Guide

Chapter 3: Matter & Energy

Classification of Matter

Matter can be classified based on its composition and properties. Understanding these classifications is fundamental in chemistry.

• Pure Substances: Have a fixed composition and distinct properties. Examples: elements (e.g., O2, Fe) and compounds (e.g., H2O, NaCl).

• Mixtures: Physical combinations of two or more substances. Can be separated by physical means.

• Homogeneous Mixtures (Solutions): Uniform composition throughout (e.g., saltwater).

• Heterogeneous Mixtures: Non-uniform composition (e.g., salad, sand in water).

Physical and Chemical Changes

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

• Chemical Change: Produces new substances with different properties (e.g., rusting iron).

States of Matter

• Solids: Definite shape and volume; particles are closely packed.

• Liquids: Definite volume, no definite shape; particles are less tightly packed than solids.

• Gases: No definite shape or volume; particles are far apart and move freely.

Physical Properties: Color, density, melting point, boiling point. Chemical Properties: Reactivity, flammability, acidity.

Chapter 4: Atoms & Elements

Periodic Table Components

• Periods: Horizontal rows.

• Groups: Vertical columns; elements in the same group have similar properties.

• Metals: Left and center; good conductors, malleable.

• Nonmetals: Right side; poor conductors, brittle.

• Metalloids: Border between metals and nonmetals; have mixed properties.

Atomic Structure

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

• Mass Number (A): Total number of protons and neutrons.

• Atomic Mass: Weighted average mass of all isotopes.

Calculating Subatomic Particles:

• Number of protons = atomic number (Z)

• Number of neutrons = mass number (A) - atomic number (Z)

• Number of electrons = number of protons (for neutral atoms)

Chapter 6: Ionic & Molecular Compounds

Naming Compounds

• Ionic Compounds: Metal + nonmetal; name the metal, then the nonmetal with "-ide" ending (e.g., NaCl = sodium chloride).

• Molecular Compounds: Nonmetal + nonmetal; use prefixes (mono-, di-, tri-, etc.) (e.g., CO2 = carbon dioxide).

Ionic Charges

• Group 1: +1, Group 2: +2, Group 17: -1, Group 16: -2, etc.

Chemical Bonds

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

• Covalent Bonds: Sharing of electrons between nonmetals.

Lewis Structures

• Show valence electrons as dots.

• Single, double, triple bonds represent shared electron pairs.

Molecular Shape (VSEPR Theory)

• Linear: 2 electron groups (e.g., CO2).

• Trigonal Planar: 3 electron groups (e.g., BF3).

• Bent: 2 bonds + 1 or 2 lone pairs (e.g., H2O).

• Tetrahedral: 4 electron groups (e.g., CH4).

• Trigonal Pyramidal: 3 bonds + 1 lone pair (e.g., NH3).

Chapter 7: Chemical Quantities & Reactions

Balancing Chemical Equations

• Ensure the same number of each atom on both sides of the equation.

Types of Chemical Reactions

• Combination (Synthesis): $A + B \rightarrow AB$

• Decomposition: $AB \rightarrow A + B$

• Single Replacement: $A + BC \rightarrow AC + B$

• Double Replacement: $AB + CD \rightarrow AD + CB$

• Combustion: Hydrocarbon + O2 $\rightarrow$ CO2 + H2O

Molar Mass and Conversions

• Molar Mass: Sum of atomic masses in a compound (g/mol).

• Conversions: Use molar mass to convert between grams and moles.

Example: To convert grams to moles: $\text{moles} = \frac{\text{grams}}{\text{molar mass}}$

Stoichiometry

• Use mole ratios from balanced equations to calculate amounts of reactants/products.

Redox Reactions

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

• Identify which species is oxidized/reduced in a reaction.

Chapter 8: Gases

Pressure Units

• 1 atmosphere (atm) = 760 mmHg = 760 Torr

Gas Laws

• Boyle’s Law: $P_1V_1 = P_2V_2$ (at constant T, n)

• Charles’ Law: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ (at constant P, n)

• Gay-Lussac’s Law: $\frac{P_1}{T_1} = \frac{P_2}{T_2}$ (at constant V, n)

• Combined Gas Law: $\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$

Chapter 9: Solutions

Concentration Calculations

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

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

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

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

Dilution Calculations

• Use: $C_1V_1 = C_2V_2$

Chapter 10: Acids & Bases

Characteristics and Naming

• Acids: Sour taste, turn litmus red, react with metals.

• Bases: Bitter taste, slippery, turn litmus blue.

• Naming: Acids: "hydro-...-ic acid" for binary acids (e.g., HCl = hydrochloric acid); Bases: named as hydroxides (e.g., NaOH = sodium hydroxide).

Bronsted-Lowry Theory

• Acid: Proton (H+) donor.

• Base: Proton acceptor.

Acid-Base Reactions

• Acid + Base $\rightarrow$ Salt + Water (neutralization)

• Identify conjugate acid-base pairs in reactions.

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, NH3).

Acid-Base Equilibrium and Le Chatelier’s Principle

• Equilibrium: Forward and reverse reactions occur at the same rate.

• Le Chatelier’s Principle: System shifts to counteract changes in concentration, temperature, or pressure.

Water as Amphiprotic

• Water can act as both an acid and a base.

• Conjugate acid: H3O+; Conjugate base: OH-

Neutralization Reactions

• Acid + Base $\rightarrow$ Salt + Water

Water Dissociation and pH Calculations

• Water dissociation: $H_2O + H_2O \rightleftharpoons H_3O^+ + OH^-$

• Ion product constant: $K_w = [H_3O^+][OH^-] = 1.0 \times 10^{-14}$ at 25°C

• pH: $pH = -\log[H_3O^+]$

• [H3O+]: $[H_3O^+] = 10^{-pH}$

pH Scale and Health

• pH < 7: Acidic; pH = 7: Neutral; pH > 7: Basic

• Blood pH is tightly regulated (7.35–7.45)

Buffers

• Solutions that resist changes in pH when small amounts of acid or base are added.

• Contain a weak acid and its conjugate base (or vice versa).