1. Chapter 19 Acids, Bases, and Salts Anything in black letters = write it in your notes (‘knowts’) UNIT 6 – Acids & Bases & Redox Rxns Chapter 19 – Acids, Bases, and Salts Chapter 20 – Oxidation-Reduction Rxns

2. 19.1 – Acid-Base Theories Taste sour Dissolve active metals to produce hydrogen gas Turns litmus paper RED Bases Taste bitter Feels slippery on skin (dissolves oils on skin) Turns litmus paper BLUE Have you seen the litmus paper yet?? Acids These are experimental definitions, they do not explain (theory) how an acid is different from a base.

3. Svante Arrhenius (1857 – 1927) Arrhenius defined an acid and base theoretically.

4. Dissociate - to split or separate from another First, a vocal word…

5. ACID – substance that dissociates in water to form hydrogen ions (H + ). BASE – substance that dissociates in water to form hydroxide ions (OH - ). HCl (aq)  H + (aq) + Cl - (aq) NaOH (aq)  Na + (aq) + OH - (aq) Arrhenius Definition (~1887)

6. When an acid is placed in water, H + ions are produced. Hydrogen ions can also be thought of as H 3 O + ions. H 3 O + = hydronium ion HCl (aq)  H + (aq) + Cl - (aq) HCl + H 2 O  H 3 O + + Cl - or equivalently,

7. Brønsted-Lowry Definition (~1923) ACID – donates H + BASE – accepts H + B-L definition covers more examples than the Arrhenius definition. Johannes Bronstad (1879 – 1947) Thomas Lowry (1874 – 1936)

8. ammoniaammonium ion water donates a H+ and so is a B-L acid ammonia accepts a H + and so is a B-L base

9. Conjugate Acid – formed when a base accepts a H + Conjugate Base – formed when an acid donates a H +

10. Amphoteric – substance that can be an acid or a base – depending on what it reacts with. Water is amphoteric

11. ACIDS donate H +, BASES accept H + label each reactant as an acid and base, label the products as conjugate acids or conjugate bases. HNO 3 + H 2 O  H 3 O + + NO 3 - CH 3 COOH + H 2 O  H 3 O + + CH 3 COO - NH 3 + H 2 O  NH 4 + + OH - H 2 O + CH 3 COO -  CH 3 COOH + OH -

12. Lewis Acids and Bases (not covered)

13. Assignment: Chapter 19 #52-57 (p. 684)

14. Acid/Base Indicators Litmus Acid – red, Base – blue, Neutral - colorless Phenolphthalein Acid – colorless, Base – pink, Neutral - colorless Cabbage Acid – red/pink, Base – yellow/green, Neutral – blue/purple

15. 19.2 – Hydrogen Ions and Acidity Molarity (M) – unit used to express the concentration of a solution Molarity = mol solute (mol) liters of soln (L) [H + ] = ‘the hydrogen ion concentration’ anything in [brackets] means the concentration in molarity [OH - ] = ‘the hydroxide ion concentration’

16. Self-Ionization of Water Water ionizes to produce a small amount of H + and OH - ions. H 2 O  H + + OH - [H + ] = [OH-] = 1 x 10 -7 M In pure water at 25 ̊ C

17. Ion-product constant for water (K w ) K w = [H + ][OH - ] = 1.0 x 10 -14 remember…anything in [brackets] represents the concentration in molarity A solution is acidic if [H + ] > 1.0 x 10 -7 M …or if the pH of the solution is below 7

18. pH = ‘power of the hydrogen ion’ pH = -log[H + ] Just as the mole was used to simplify large numbers of atoms, pH is used to simplify small concentration numbers [H + ] = 1 x 10 -7 M Instead of writing out numbers like these… [H + ] = 2.4 x 10 -4 M [H + ] = 7.3 x 10 -10 M pH = 7.00 pH = 3.62 pH = 9.14 we can write number like these

19. We can just say “This solution has a pH of 3.62”. Instead of saying “This solution has a hydronium ion concentration of 2.4 x 10 -4 M”. Not only is pH an easier number to talk about, pH is understood by most people, whereas molarity is not. The pH scale is used to describe how acidic or basic (alkaline) a substance is.

21. Pure water has [H + ] = 1.00 x 10 -7 M pH = - log [1.00 x 10 -7 ] pH = 7 The pH of water would be pH = -log[H + ] Examples

22. [H + ] = 2.3 x 10 -5 M. Calculate the pH. pH = - log [H + ] pH = - log [2.3 x 10 -5 ] pH = 4.64

23. Examples [H + ] = 1.0 x 10 -5 M. Calculate the pH. pH = - log [H + ] pH = - log [1.0 x 10 -5 ] pH = 5.0

24. Examples -4.2 = log [H + ] 10 -4.2 = 6.31 x 10 -5 M = [H + ] pH = - log [H + ] 10 -4.2 = 10 log [H+] pH = 4.2. Calculate [H + ]

25. Summary of pH [H + ] = 10 -pH pH = - log [H + ] The pH scale is used to indicate how acidic or basic a substance is. The scale normally ranges from 0 to 14. Acids have a pH below 7, bases are above 7.

26. ASSIGNMENT: Chapter 19 #10-24 (p.655-662)

27. 19.3 – Strengths of Acids and Bases Coming soon!!

28. 19.4 – Neutralization Reactions Neutralization Rxn – complete rxn of a strong base with a strong acid A neutralization rxn will produce a salt and water. Acid + Base  Salt + H 2 O HCl + NaOH  NaCl + H 2 O

29. Titration – determining the concentration of an unknown solution using a solution whose concentration is known. Equivalence Point – point where the amount of acid equals the amount of base Standard – solution of known concentration. End Point – point where the indicator changes color

30. EXAMPLE 10.0 mL of 0.5 M HCl solution is added to 20.0 mL of NaOH of unknown concentration. What is the concentration of the NaOH? HCl + NaOH  NaCl + H 2 O 0.5 M 10.0 mL20.0 mL x M Since the reaction of HCl and NaOH is 1:1 and twice the volume of NaOH was used, the NaOH must half as strong as HCl; [0.25 M].

31. EXAMPLE What volume of 0.10 M KOH is required to neutralize 20.0 mL of 0.20 M H 2 SO 4 solution? H 2 SO 4 + 2KOH  K 2 SO 4 + 2H 2 O 0.20 M 20.0 mLx mL 0.10 M Always adjust if the rxn is NOT 1:1 Since KOH requires twice as many moles as H 2 SO 4, you should double your answer.

32. 19.5 – Salts in Solution not covered…

33. Neutralization Rxn – Acid + Base  Salt + H 2 O HCl + NaOH  NaCl + H 2 O

34. Sulfuric acid + magnesium hydroxide Write the balanced chemical equation for each neutralization reaction Phosphoric acid + calcium hydroxide Nitric acid + ammonium hydroxide

35. Chapter 19 Quiz Questions 1. What color will litmus paper be in an acidic solution? 3. Define an acid in terms of the Arrhenius definition. 4. What two products are always formed in an acid-base neutralization reaction? 2. What color will phenolphthalein indicator be in an basic solution?

36. 5. A student titrated 10.0 mL of an HCl solution. The titration required 23.3 mL of 0.24M NaOH solution. a.Which solution was the standard? b.Which solution was more concentrated? c.Convert the volumes to liters d.Calculate the number of moles of NaOH that reacted. e.Calculate the number of moles of HCl that reacted. f.Calculate the molarity of the HCl solution.

37. 6. Calculate the pH of solutions with the following hydrogen ion concentrations. a. [H + ] = 1.23 x 10 -4 M b. [H + ] = 3.42 x 10 -7 M 7. Calculate the hydrogen ion concentrations of solutions with the given pH. a. pH = 3.14 b. pH = 9.2