Reactions in Aqueous Solutions II: Calculations

1. Reactions in Aqueous Solutions II: Calculations Chapter 11

2. Calculations Involving Molarity • Review sections on molarity and percent by mass (Chapter 3) • If 100.0 mL of 1.00 M NaOH solution and 100.0 mL of 0.500 M H2SO4 solution are mixed, what will be the resulting molarity of the salt? Assume that the amount of water produced the reaction is negligible. What if it’s not? Does anything else remain? • Volumes are additive

3. Calculations Involving Molarity • If 130 mL of 1.00 M KOH and 100 mL of 0.500 M H2SO4 solutions are mixed, what will be the concentration of the products in solution? • You need to remember how to determine the limiting reaction. This will determine the amount of product formed. From the amount of limiting reactant, the amount of excess reactant can be determined. • In a flask, 1.29 grams of Mg(OH)2 is added to 50.0 mL of water. This is reacted with 35.0 mL of a 1.50 M HCl solution, what will be the concentration of the products in the solution? Note: Write the reactions for the formation of the normal salts.

4. Calculations Involving Molarity • What volume of 0.750 M NaOH solution would be required to neutralize completely 100 mL of 0.250 M H3PO4? • This is a complete neutralization. • What type of reaction is this? What will be the products and direction of the arrows? There are multiple ways to solve these problems. Always start with the balanced equation.

5. Titrations • Titration - process in which a solution of one reactant, the titrant, is carefully added to a solution of another reactant, and the volume of titrant required for complete reaction is measured. • Equivalence point – The point at which stoichiometrically equivalent amounts of an acid and base have reacted. • Indicator – A compound that exhibits its different colors in solutions of different acidities. Used to determine the point at which an acid-base reaction is complete. • End point – The point at which an indicator changes color and a titration is stopped. The end point should coincide with the equivalence point. DEMO: Titration with an indicator. Use the numbers for a future problem.

6. Titrations • What is the molarity of a freshly-made HCl solution if 36.7 mL of the HCl solution is neutralized with 28.2 mL of 0.355 M NaOH solution? • Previous titration. • What is the molarity of a calcium hydroxide solution if 41.9 mL of the solution is neutralized with 19.1 mL of 0.418 M HCl solution?

7. Titrations • Standard solutions – solutions of accurately known concentrations. • Primary standard – A substance of a known high degree of purity that undergoes one invariable reaction with the other reactant. • Look at page 408 for the properties of a primary standard • Secondary standard – A solution that has been titrated against a primary standard. • Standardization – The process by which the concentration of a solution is accurately determined by titrating with a primary standard.

8. Using Primary Standards for Titrations • A commonly primary standard for solutions of acids is Na2CO3. Even though this is a ____ salt, it reacts with acids. • H2SO4 + Na2CO3 Na2SO4 + CO2 + H2O • How does the reaction produce CO2 and H2O? • Calculate the molarity of a sulfuric acid solution if 23.2 mL of it reacts with 0.212 g of Na2CO3. • Calculate the molarity of a HCl solution if 28.5 mL of it reacts with 0.458 grams of Na2CO3.

9. Using Primary Standards for Titrations • A common primary standard for solutions of bases is the acidic salt, potassium hydrogen phthalate, KC6H4(COO)(COOH) or KHP. • This acidic salt is obtained by reacting phthalic acid with KOH (see page 409). One hydrogen remains to neutralize bases. • The salt is highly soluble and can be obtained in a high state of purity. • Molecular weight is ____

10. Using Primary Standards for Titrations • Calculate the molarity of a NaOH solution if 27.3 mL of it reacts with 0.4084 g of KHP. • What if Sr(OH)2 reacts with KHP? Write the balanced equation. • An impure sample of oxalic acid, (COOH)2, had a mass of 0.1743 g. It was dissolved in water and titrated with 39.82 mL of 0.08915 M NaOH solution. Calculate the percent purity of the (COOH)2sample.

11. Equivalent Weights and Normality • Normality is the number of equivalent weights of solute per liter of solution. • Equivalent weight of an acid is the mass in grams necessary to furnish NA of H+ ions. • Monoprotic (HCl) 1 mol = 1 eq • Diprotic (H2SO4) 1 mol = 2 eq • Triprotic (H3PO4) 1 mole = 3 eq

12. Equivalent Weights and Normality • Works the same way for bases • NaOH, Ca(OH)2, and Al(OH)3 • Table 11-1 in textbooks • What is the mass of one equivalent of H2SO4? • What is the mass of one equivalent of Al(OH)3? • Calculate the normality of a solution that contains 196 g of sulfuric acid in 1500 mL of solution.

13. Equivalent Weights and Normality • Calculate the molarity and normality of a solution that contains 34.2 g of barium hydroxide in 8.00 liters of solution. • One equivalent of an acid reacts with one equivalent of a base. It follows that for an acid-base reaction that goes to completion. Number of equivalents of acid = Number of equivalents of base

14. Equivalent Weights and Normality • The product of the volume and its normality is equal to the number of equivalents of solute contained in the solution. • This relationship also works for a base. • What volume of 6.0 M phosphoric acid solution is required to prepare 900 mL of 0.200 N phosphoric acid solution?

15. Equivalent Weights and Normality • What is the normality of a sulfuric acid solution if 31.3 mL of it reacts with 0.318 g of sodium carbonate? • 30.0 mL of 0.750 N nitric acid solution required 22.5 mL of calcium hydroxide solution for neutralization. Calculate the normality and the molarity of the calcium hydroxide solution.