2. I. Introduction to Acids & Bases Acids & Bases

3. Properties

4. B. Definitions  Arrhenius - In aqueous solution… HCl + H 2 O  H 3 O + + Cl – AcidsAcids form hydronium ions (H 3 O + ) H HHHH H Cl OO – + acid

5. B. Definitions  Arrhenius - In aqueous solution… BasesBases form hydroxide ions (OH - ) NaOH → Na + + OH -

6. B. Definitions  Brønsted-Lowry HCl + H 2 O  Cl – + H 3 O + AcidsAcids are proton (H + ) donors. BasesBases are proton (H + ) acceptors. conjugate acid conjugate base baseacid

7. B. Definitions H 2 O + HNO 3  H 3 O + + NO 3 – CBCAAB The conjugate base forms after the acid donates a Hydrogen The conjugate acid forms after the base accepts a Hydrogen

8. B. Definitions - can be an acid or a base.  Amphoteric - can be an acid or a base. NH 3 + H 2 O  NH 4 + + OH - CACBBA

9. B. Definitions F - H 2 PO 4 - H2OH2O HF H 3 PO 4 H 3 O +  Give the conjugate base for each of the following: - an acid with more than one H +  Polyprotic - an acid with more than one H +

10. B. Definitions Br - HSO 4 - CO 3 2- HBr H 2 SO 4 HCO 3 -  Give the conjugate acid for each of the following:

11. C. Neutralization Reactions  Acids react with bases to produce a salt and water NaOH (aq) + HCl(aq) → NaCl + H 2 O base acid salt water Ca(OH) 2 + 2 HCl → CaCl 2 + 2H 2 O base acid salt water

12. C. Neutralization Salt  Acid + Base  Salt + Water  HCl + NaOH  NaCl + HOH salt water

13. Acids  Acids can be recognized because the start with H  Examples HCl H 2 SO 4 HI

14. Acids  Acids are in aqueous solution (aq)  For the purposes of this class, we will assume that if it begins with H, we will name it according to the rules of naming acids

15. Rule #1 - naming acids  If the anion ends in –ide, the acid will be named…  Hydro (root) – ic acid  This is usually for H plus one element

16. For example  HCl  Hydrochloric acid  HI  Hydroiodic acid H2SH2S  Hydrosulfuric acid

17. Rule #2 – naming acids  If you have an H plus an anion ending in –ate, the acid will be named…  (root) – ic acid

18. Examples  H 2 SO 4  Sulfuric acid  HNO 3  Nitric acid  H 3 PO 4  Phosphoric acid

19. Rule # 3 – naming acids  If you have an H plus an anion ending in –ite, the acid will be named…  (root) – ous acid

20. Examples  H 2 SO 3  Sulfurous acid  HNO 2  Nitrous acid  H 3 PO 3  Phosphorous acid

21. Remember… ate  ic ite - ous

22. Writing formulas for acids  When writing formulas for acids you MUST look at the charges and bring them down!

23. Examples  HBr  Hydrogen + one element  Hydrobromic acid  HClO 3  H + chlorate  ate  ic  Chloric acid

24. More examples  H 2 SO 3  H 2 CO 3  HF  Nitrous acid  Perchloric acid  Iodic acid  Sulfurous acid  Carbonic acid  Hydrofluoric acid  HNO 2  HClO 4  HIO 3

25. Acids & Bases II. pH

26. A. Ionization of Water H 2 O + H 2 O H 3 O + + OH - K w = [H 3 O + ][OH - ] = 1.0  10 -14 [ ] = Concentration (Molarity)

27. A. Ionization of Water  Find the hydroxide ion concentration of 3.0  10 -2 M HCl. [H 3 O + ][OH - ] = 1.0  10 -14 [3.0  10 -2 ][OH - ] = 1.0  10 -14 [OH - ] = 3.3  10 -13 M Acidic or basic? Acidic

28. pH = -log[H 3 O + ] B. pH Scale 0 7 INCREASING ACIDITY NEUTRAL INCREASING BASICITY 14 pouvoir hydrogène (Fr.) “hydrogen power”

29. B. pH Scale pH of Common Substances

30. B. pH Scale pH = -log[H 3 O + ] pOH = -log[OH - ] pH + pOH = 14 [H 3 O + ] = Inverse log (-pH) [OH - ] = Inverse log (-pOH)

31. B. pH Scale  What is the pH of 0.050 M HNO 3 ? pH = -log[H 3 O + ] pH = -log[0.050] pH = 1.3 Acidic or basic? Acidic

32. B. pH Scale  What is the molarity of HBr in a solution that has a pOH of 9.6? pH + pOH = 14 pH + 9.6 = 14 pH = 4.4 Acidic pH = -log[H 3 O + ] 4.4 = -log[H 3 O + ] -4.4 = log[H 3 O + ] Inverse log (–pH) = Molarity [H 3 O + ] = 4.0  10 -5 M HBr