1. Ch 9: Acids, Bases and Salts Suggested Problems: 2, 6, 10, 12, 28-44, 82, 94-100, Bonus: 118

2. Acids and Bases in Aqueous Solution Think back to Chapter 4 –Acid Definition: A substance that produces H + when dissolved in H 2 O Examples: HCl, HNO 3, H 2 SO 4 –Base Definition: A substance that produces OH - when dissolved in H 2 O –Examples: KOH, NaOH, NH 3

3. Common Acids Sulfuric Acid H 2 SO 4 Phosphoric acid H 3 PO 4 Acetic acid CH 3 CO 2 H Nitric acid HNO 3 Hydrochloric acid HCl

4. Common Bases Sodium Hydroxide NaOH Calcium hydroxide Ca(OH) 2 Magnesium hydroxide Mg(OH) 2 Ammonia NH 3

5. Common Acids and Bases Acids –H 2 SO 4 –HCl –H 3 PO 4 –HNO 3 –CH 3 CO 2 H Chemical Formulas Begin with H Formulas Contain CO 2 H Bases –NaOH –Ca(OH) 2 –Mg(OH) 2 –NH 3 Chemical Formulas Contain OH

6. Arrhenius Acids and Bases 1903 Chemistry Nobel Prize –Barely Awarded Ph.D. Technicality issue with Arrhenius acid definition –H + is very reactive

7. Updated Definitions Arrhenius acids –Substances that produce H 3 O + when dissolved in H 2 O Arrhenius bases –Substances that produce OH - when dissolved in H 2 O What if the reactions are not in H 2 O?

8. Brønsted-Lowry Acids and Bases Separately developed the same theory pertaining to acids and bases in 1923 Johannes Brønsted Thomas Lowry

9. Brønsted-Lowry Acid Definition: any substance that is able to give a hydrogen ion H +, to another molecule or ion –Proton donor Not limited to reactions in H 2 O –Do not have to create appreciable [H 3 O + ] –NaOH (s) + HCl (aq) CH 3 CO 2 H (aq) + H 2 O (l) H 3 O + (aq) + CH 3 CO 2 - (aq) NaCl (aq) + H 2 O (l)

10. Brønsted-Lowry Acids Different acids can donate different numbers of H + Acid# of Acidic H + Terminology HCl H 2 SO 4 H 3 PO 4 CH 3 CO 2 H 1 2 3 1 monoprotic diprotic triprotic monoprotic

11. Brønsted-Lowry Bases Definition: a substance that accepts H + from an acid –Proton Acceptor Not limited to reactions in H 2 O –Do not have to create appreciable [OH - ] –NH 3(g) + HCl (g) CH 3 CO 2 H (aq) + H 2 O (l) H 3 O + (aq) + CH 3 CO 2 - (aq) B-L Base NH 4 Cl (s)

12. Identify the following as Brønsted-Lowry Acids, Bases or Neither HCN AlCl 3 H 2 CO 3 CH 3 CO 2 - Mg 2+ CH 3 NH 3 + Acid Base Neither Acid Neither

13. Brønsted-Lowry Acids and Bases Summary –Acid-Base reaction is one in which a proton is transferred Use B or B - to represent bases Use HA to represent acids Using the symbols B, B - and HA write two general acid-base chemical reactions B + HA BH + + A - B - + HA BH + A -

14. Brønsted-Lowry Acids and Bases Consequence of Brønsted-Lowry Definition –What are species BH +, BH, and A - ? –BH + –BH –A-–A- –Acid-Base reactions are reversible K is often large, resulting in a single preferred direction Acid Base

15. Conjugate Acid – Base Pairs Definition: A pair of compounds whose formula differ only by one proton. –After an acid donates a proton, the remaining species turns into a conjugate base (CB). –After a base accepts a proton, the resulting species turns into a conjugate acid (CA).

16. HF (aq) + H 2 O (l) H 3 O + (aq) + F - (aq) - H + +H + AcidBaseConjugate Acid Conjugate Base + H + Conjugate Acid – Base Pairs Example

17. NH 3(g) + H 2 O (l) NH 4 + (aq) + OH - (aq) + H + BaseAcidConjugate Acid Conjugate Base - H + Conjugate Acid – Base Pairs Example

18. HF (aq) + H 2 O (l) H 3 O + (aq) + F - (aq) Water as Both an Acid and a Base NH 3(g) + H 2 O (l) NH 4 + (aq) + OH - (aq) A substance that can react as an acid or a base is called amphoteric Acid Base

19. Common Acid-Base Reactions Neutralization reaction –Acid with a metal hydroxide Salt: anion of the acid with the cation of the base –HCl (aq) + KOH (aq) KCl (aq) + H 2 O (l) –Why is this called a neutralization reaction? Net Ionic Equation –H + (aq) + OH - (aq) H 2 O (l) Products are always salt and H 2 O

20. Acid-Base Reactions Write the balanced chemical equation for the reaction of sulfuric acid with magnesium hydroxide.

21. Acid-Base Reactions Acid with bicarbonate and carbonate ions –Bicarbonate ion HCO 3 - –H + (aq) + HCO 3 - (aq) [H 2 CO 3(aq) ] CO 2(g) + H 2 O (l) –Carbonate ion CO 3 2- –2H + (aq) + CO 3 2- (aq) [H 2 CO 3(aq) ] CO 2(g) + H 2 O (l) Products of these reactions are salt, CO 2 and H 2 O

22. Acid-Base Reactions Write the balanced chemical reaction for nitric acid with baking soda (sodium bicarbonate).

23. Acid-Base Reactions Acid with Ammonia –Products for this general reaction are ammonium salts NH 3(aq) + HNO 3(aq) NH 4 NO 3(aq) Write the balanced chemical reaction for ammonia with sulfuric acid.

24. Challenge Problem An over the counter antacid has NaAl(OH) 2 CO 3 as the active ingredient. –How many grams of this antacid are required to nuetralize 15.0 mL of 0.0955 M HCl?

25. The Self Ionization of Water H 2 O is amphoteric But what if you have just pure water? H 2 O (l) + H 2 O (l) ⇌ H 3 O + (aq) + OH - (aq) This equilibrium is governed by the equilibrium constant K w

26. Equilibrium Constant Equilibrium constant (K) is equal to the concentration of the products divided by the reactants aA + bB cC + dD [x] = concentration of species X in molarity

27. KWKW H 2 O (l) + H 2 O (l) ⇌ H 3 O + (aq) + OH - (aq)

28. Strong Acids Strong acids give away all of their hydrogen ions For example, HCl is a strong acid, and when HCl dissolves in water: HCl  H + + Cl -

29. Weak Acids Weak acids do not give away their H + ions, and are in equilibrium with their ionized form Most acids are weak acids For example, acetic acid is a weak acid, and when HC 2 H 3 O 2 dissolves in water: HC 2 H 3 O 2 ⇌ H + + C 2 H 3 O 2 -

30. Strong Bases A strong base will give away all of its hydroxide ions (OH - ) For example, NaOH is a strong base, and when NaOH dissolves in water: NaOH  Na + + OH -

31. Weak Bases To think about weak bases you must think in terms of a proton acceptor not in terms of OH -. ( Brønsted-Lowry Base) Weak bases accept some H +. Again as with weak acids there is an equilibrium present.

32. Strong/Weak Acids and Bases The description of strong/weak has nothing to do with concentration Concentration is independent of it being strong or weak. Concentration is a measure of the amount of moles per liter You can have low concentrations of Strong Acid and Bases

33. pH and pOH pH informs a person about whehter or not a solution is acid or basic pH = -log[H + ] pOH = -log[OH - ] pH of 7 is nuetral pH less than 7 is acidic pH greater than 7 is basic pH + pOH = 14

34. pH and pOH calculation examples Determine the concentration of H 3 O + and OH - from the following pH values pH = 9.0 pH = 3.0 pH = 11.0

35. pH and pOH Calculations This course will deal only with non- equilibrium acids and bases when calculating pH or pOH Therefore the concentration of H + and OH - will be able to be determined from the stoichiometry of the formula. For example –What is the pH of a solution of 0.10 M HCl? –What is the pH of a solution of 0.20 M NaOH?

36. Salts Definition: a substance composed of the cation of a base with the anion of the acid Need to discuss Equivalent units –This term is terribly misused by the medical and biological profession A equivalent is the quantity of material necessary to deliver one unit of chemical reactivity –It makes no sense outside of the context of a chemical reaction! However, in blood analysis, Equivalents = moles x charge on ion

37. Equivalents Example Determine the number of equivalents in the following: 0.10 mol of NaCl 0.10 mol of CaCl 2 –Only consider either the positive or negative charges not both and the origination of the species is also important

38. Equivalent Example A sample of blood serum contains 0.139 eq/L of Na + ion. Assume the Na + comes from dissolved NaCl, and calculate the number of equivalents, number of moles and number of grams of NaCl in 250 mL of the serum.

39. Titration Calculation A 25 mL sample of vinegar (which contains acetic acid) is titrated with 0.100 M NaOH. If 6.75 mL of NaOH are required, what is the molarity of the acetic acid in vinegar? 25 mL of vinegar 0.100 M NaOH

40. Titration Example A 25.0 mL sample of H 2 SO 4 solution requires the addition of 16.3 mL of 0.200 M NaOH solution to reach the equivalence point. What is the concentration of the acid?

41. Buffers A pH buffer is a solution that resists changes in pH A pH buffer must contain a weak acid (HA) and its conjugate base (A - ) HA + OH -  H 2 O + A - A - + H +  HA

42. Buffer Example Carbonic acid and bicarbonate are important blood buffers