Types of Chemical Reactions in Aqueous Solutions: Precipitation, Acid-Base, and Oxidation-Reduction Reactions

1. Chapter 4 Types of Chemical Reactions and Solution Stoichiometry

2. Section 4.1 Water, The Common Solvent • Nature of Water • Shape • Chemical make-up • Polar molecule – due to the partial charge on the molecule due to the large oxygen attraction for electrons. • Hydration • Solubility

3. Section 4.2 The Nature of Aqueous Solutions • Solute • Solvent • Electrical Conductivity • Strong electrolytes • Weak electrolytes • Nonelectrolytes

4. Section 4.3 The Composition of Solutions • Molarity – defined as moles of solute per volume of solution in liters • Calculations for Molarity pages 141 – 144 • Standard solution – a solution whose concentration is accurately known. Dilution • Stock solution • M1V1 = M2V2 • Calculations for dilution – pages 145 - 148 • Pipets and Volumetric flasks • Volumetric and measuring pipets

5. Section 4.4 Types of Chemical Reactions • Types of Solution Reactions • Precipitation reactions • Acid-base Reactions • Oxidation – reduction reactions

6. Precipitation reactions – when two solutions are mixed, an insoluble substance forms. Here the solid “drops” or precipitates out of solution Precipitate Calculations for reactions on pages 153 - 154 Rules for the solubility of Salts in Water Most nitrate salts are soluble Most salts containing the alkali metal ions and the ammonium ion are soluble. Most chloride, bromide and iodide salts are soluble. Notable exceptions are salts containing the ions Ag+, Pb+2, and Hg+2. Most sulfate salts are soluble. Excepts are BaSO4, PbSO4, HgSO4, and CaSO4. Most hydroxide salts are only slightly soluble. The important soluble hydroxides are NaOH and KOH. The compounds Ba(OH)2, Sr(OH)2, and Ca(OH)2 are marginally soluble. Most sulfide (S-2), carbonate(CO3-2), chromate(CrO4-2), and phosphate (PO4-3) are only slightly soluble Section 4.5 Precipitation Reactions

7. Section 4.6 Describing Reactions in Solution • Types of Equations • Molecular equation • Complete Ionic equation • Spectator ions • Net ionic equation • Practice for writing equations pages 155 - 156

8. Section 4.7 Stoichiometry of Precipitation Reactions • Steps for solving Precipitation Reaction problems • Identify the species present in the combined solution, and determine what reaction occurs. • Write the balanced net ionic equation for the reaction. • Calculate the moles of reactants. • Determine which reactant is limiting. • Calculate the moles of product or products, as required. • Convert to grams or other units, as required. • Practice problems – pages 156 - 158

9. Acids – proton donor Bases – proton acceptor The hydroxide ion is such a strong base that for purposes of stoichiometric calculations it can be assumed to react completely with any weak acid that this class will encounter. Practice problems on pages 159 - 161 Steps for Calculating Acid-Base Reactions List the species present in the combined solution before any reaction occurs, and decide what reaction will occur. Write the balanced net ionic equations for the reaction. Calculate the moles of reactants. For reactions in solution, use the volumes of the original solutions and their molarities. Determine the limiting reactant where appropriate. Calculate the moles of the required reactant or product. Convert to grams or volume (of solution), as required. Section 4.8 Acid-Base Reactions

10. Section 4.8 Acid/Base Titrations • Volumetric analysis • Titration • Equivalence point • Stoichiometric point • Indictator • Endpoint • Three requirements for a successful titration: • The exact reaction between titrant and analyte must be known (and rapid). • The stoichiometric (equivalence) point must be marked accurately. • The volume of titrant required to reach the stoichiometric point must be known accurately. • Practice for Acid/base titrations (neutralization) pages 162 - 164

11. Section 4.9 Oxidation/Reduction Reactions • Oxidation – reduction reactions – aka Redox reactions are those in which one or more electrons are transferred. • Oxidation states • Rules for Assigning Oxidation States • An atom in a neutral state is zero • A monatomic ion is the same as its charge • Fluorine is -1 in its compounds • Oxygen is -2 in its compounds, with the exception of peroxides O2-2 in which oxygen is -1 • Hydrogen is +1 in its covalent compounds • Characteristics • Oxidation • Reduction • Oxidizing agent • Reducing agent • Practice Problems pages 170 - 171

12. Section 4.10 Balancing Oxidation – Reduction Equations • Half-reaction • Half-reaction Method for Balancing Equations for Redox Reactions in Acidic Solutions • Write separate equations for the oxidation and reduction half-reactions • For each half-reaction: • Balance all the elements except hydrogen and oxygen • Balance oxygen by using water • Balance hydrogen using H+ • Balance the charge using electrons • If necessary, multiply one or both balanced half-reactions by an integer to equalize the number of electrons transferred in the two half-reactions. • Add the half-reactions, and cancel identical specials. • Check that the elements and charges are balanced. • Practice problems page 173 - 176

13. Balancing Redox Continued… • The Half-reaction method for balancing equations for redox reaction in basic solutions • Use the half-reaction method as specified for acidic solutions to obtain the final balanced equations as if H+ ions were present. • To both sides of the equation obtained above, add a number of OH- ions that is equal to the number of H+ ions.( We want to eliminate H+ by forming water) • Form water on the side contain both H+ and OH- ions, and eliminate the number of water molecules that appear on both sides of the equation. • Check that elements and charges are balanced. • Practice for balancing this type of reaction - pages 178 - 179