Unit 4 Review: Chemical Reactions - Key Concepts

Physical vs. Chemical Changes

• Physical change: alters form but not chemical identity; phase changes involve intermolecular forces, not bond breaking/formation.
• Chemical change: creates new substances via bond breaking and formation; also known as a chemical reaction.
• Signs of chemical reactions: color change, temperature change (heat or light), gas formation, precipitate formation, or odor production.

Net Ionic Equations

• Molecular equation: balanced equation with all reactants and products.
• Complete ionic equation: separates all ions in the reaction.
• Net ionic equation: includes only participating species; excludes spectator ions.

Representations of Reactions

• Balanced chemical equations can be translated into symbolic particulate representations.

Physical and Chemical Changes

• Chemical processes involve breaking/formation of chemical bonds.
• Physical processes involve changes in intermolecular interactions (e.g., phase changes).
• Some physical processes can involve breaking chemical bonds (e.g., dissolution of salt in water).

Stoichiometry

• Atoms are conserved in chemical processes, allowing calculation of product/reactant amounts.
• Coefficients in balanced equations show proportionality of substances; used in mole concept calculations.
• Stoichiometric calculations can be combined with the ideal gas law and molarity.

Introduction to Titration

• Titration: determines analyte concentration using a titrant of known concentration that reacts specifically and quantitatively with the analyte.
• Equivalence point: analyte is totally consumed, indicated by a property change (endpoint).

Types of Chemical Reactions

• Acid-Base: involves proton transfer.
• Redox: involves electron transfer, indicated by oxidation number changes; combustion is a subclass.
• Precipitation: mixing aqueous ions to form an insoluble compound. All sodium, potassium, ammonium, and nitrate salts are soluble in water.

Rules for Oxidation Numbers

• Elements in elemental form: 0.
• Monatomic ion: same as its charge.
• Oxygen: −2, except in O_2^{2-}. Hydrogen: −1 with metal, +1 with nonmetal.
• Fluorine: always −1. Other halogens: −1 when negative.
• Sum of oxidation numbers: 0 in neutral compound, charge on ion in polyatomic ion.

Introduction to Acid-Base Reactions

• Bronsted-Lowry acid: proton donor.
• Bronsted-Lowry base: proton acceptor.
• Water can act as acid or base (amphiprotic).
• Conjugate acid: forms after base gains a proton.
• Conjugate base: remains after acid loses a proton.

Strong Acids and Bases

• Strong acids: HCl, HBr, HI, HNO3, H2SO4, HClO4, HClO_3.
• Strong bases: LiOH, NaOH, KOH, RbOH, CsOH, Sr(OH)2, Ca(OH)2, Ba(OH)_2.
• K{eq} > 1: products favored, reactants stronger; K{eq} < 1: reactants favored, products stronger.

Oxidation-Reduction (Redox) Reactions

• Oxidation: loss of electrons.
• Reduction: gain of electrons.
• OIL RIG: Oxidation Is Loss, Reduction Is Gain.