Chapter 3: Matter & Energy

Classification of Matter

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

• Pure Substances: Matter with a fixed composition; includes elements and compounds.

• Mixtures: Physical combinations of two or more substances; can be homogeneous (uniform composition, e.g., saltwater) or heterogeneous (non-uniform, e.g., salad).

• Elements: Substances made of one type of atom (e.g., O2).

• Compounds: Substances composed of two or more elements chemically combined (e.g., H2O).

Physical and Chemical Changes

• Physical Change: Alters the state or appearance without changing 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 can move past each other.

• 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.

Chapter 4: Atoms & Elements

Periodic Table Components

• Periods: Horizontal rows.

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

• Metals: Left side; shiny, conductive, malleable.

• Nonmetals: Right side; poor conductors, brittle.

• Metalloids: Along the staircase; properties of both metals and nonmetals.

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

• Number of neutrons = mass number – atomic number

• 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 (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 around symbols; lines represent bonds.

• Single, double, and triple bonds represent 2, 4, and 6 shared electrons, respectively.

Molecular Shape (VSEPR Theory)

• Linear: 180° (e.g., CO2)

• Trigonal Planar: 120° (e.g., BF3)

• Bent: <120° or <109.5° (e.g., H2O)

• Tetrahedral: 109.5° (e.g., CH4)

• Trigonal Pyramidal: ~107° (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: A + B → AB

• Decomposition: AB → A + B

• Single Replacement: A + BC → AC + B

• Double Replacement: AB + CD → AD + CB

• Combustion: Hydrocarbon + O2 → 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.

Stoichiometry

• Use coefficients in balanced equations as mole ratios to convert between substances.

Redox Reactions

• Oxidation: Loss of electrons.

• Reduction: Gain of electrons.

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

Chapter 8: Gases

Pressure Units

• 1 atm = 760 mmHg = 760 Torr

Gas Laws

• Boyle’s Law: (at constant T and n)

• Charles’ Law: (at constant P and n)

• Gay-Lussac’s Law: (at constant V and n)

• Combined Gas Law:

Chapter 9: Solutions

Concentration Calculations

• Mass Percent (m/m):

• Volume Percent (v/v):

• Mass/Volume Percent (m/v):

• Molarity (M):

Dilution Calculations

• (where C = concentration, V = volume)

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 often start with "hydro-" or end with "-ic"; bases are usually named as hydroxides (e.g., NaOH = sodium hydroxide).

Bronsted-Lowry Definitions

• Acid: Proton (H+) donor.

• Base: Proton (H+) acceptor.

Acid-Base Reactions

• Acid + Base → 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., CH3COOH, NH3).

Acid-Base Equilibrium & 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 → Salt + Water

• Write balanced equations for these reactions.

Water Dissociation and pH Calculations

• Water dissociation:

• at 25°C

• pH:

• [H3O+]:

pH Scale and Health

• pH scale: 0 (acidic) to 14 (basic); 7 is neutral.

• Blood pH is tightly regulated (7.35–7.45); deviations can be harmful.

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).