Aqueous Reactions & Solution Stoichiometry Overview

1. Unit 6 – Aqueous Reactions and Solution Stoichiometry

2. Solutions and Electrolytes • Solution is a homogeneous mixture of two or more substances • Solvent • Solute • Electrolyte form ions in aqueous solution • Ionic compounds • H2O • Non electrolyte do not form ions in solution • Molecular compounds

3. Strong vs Weak Electrolytes • Electrolyte strength determined amount of ions present in aqueous solutions • Strong Electrolytes – dissolved compound that exists mainly or completely as ions • Acids, Ionic Compounds • HCl → H+ + Cl- • Weak Electrolytes – dissolved compound that exists mainly as molecules not ions • HC2H3O2 ↔ H+ + C2H3O2-

4. Precipitation Reactions • Precipitation reactions result in formation of insoluble solid • Solubility is the amount of a substance that can be dissolved in a large amount of solvent • Compound with solubility less than 0.01mol/L is insoluble • Can not determine solubility based on physical properties

5. Solubility Rules

6. Predicting Solubility • Determine whether the following compounds are soluble or insoluble and why. • Sodium carbonate • Lead sulfate • Barium nitrate • Cobalt (II) hydroxide • Ammonium phosphate

7. Predicting Precipitation Reactions • To predict whether or not a precipitate forms look at all possible combinations of present ions and see if any form an insoluble solid. • Write the balance reactions for the following reactions. • Barium chloride and Potassium sulfate • Iron (III) sulfate and Lithium hydroxide • Barium nitrate and potassium hydroxide

8. Ionic Equations • Molecular equations • Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + 2KNO3 (aq) • Complete Ionic equations • Pb2+ + 2NO3- + 2K+ + 2I- → PbI2 + 2K+ + 2NO3- • Spectator ions • Nitrate, potassium • Net Ionic equation • Pb2+ + 2I- → PbI2

9. Writing Net Ionic Equations • Write balanced reactions • Rewrite equation to show ions in solution • Identify and cancel spectator ions • If all ions are spectator ions no reaction occurs

10. Net Ionic Equation Practice • Write the net ionic equations for the reactions between: • Calcium chloride and sodium carbonate • Silver nitrate and potassium phosphate

11. Acid and Base Reactions • Acids are substances that ionize in aqueous solution to form H+ ions • Monoprotic • HNO3 → H+ + NO3- • Diprotic • H2SO4 → 2H+ + SO42- • Bases are substances that accept H+ ions • Produce OH- ions when dissolved in water

12. Strong and Weak Acid and Bases • Strong acids, bases = Strong electrolytes • Weak acids, bases = Weak electrolytes

13. Electrolytes Summary

14. Classifying Electrolytes • Classify the following substances as strong, weak or nonelectrolytes. • Calcium chloride • Nitric Acid • Ethanol (C2H5OH) • Formic Acid (HCHO2) • Potassium hydroxide • Rank solutions of calcium nitrate, table sugar, sodium acetate, and acetic acid in order of increasing conductivity.

15. Neutralization Reactions • Acids and base can change the color of dyes • Litmus • Mix acids and bases and neutralization reaction occurs • HCl + NaOH → H2O + NaCl • H+ + OH- → H2O • Write the net ionic equation for reactions between • Hydrochloric acid and magnesium hydroxide • Acetic acid and barium hydroxide

16. Acid base Reactions • Reactions that form gases • Bases other than OH- may react with acids to form molecules • 2HCl + Na2S → H2S + 2NaCl • 2H+ + S2- → H2S • Carbonates and bicarbonates react with acids to form CO2 • HCl + NaHCO3 → NaCl + H2CO3 • H2CO3 → H2O + CO2 • HCl + NaHCO3 → NaCl + H2O + CO2 • H+ + HCO3- → H2O + CO2

17. Homework • 3, 5, 10, 11, 14, 15, 18, 23, 26, 30, 31