1. Acids and Bases

2. Properties of Acids Aqueous solution have sour taste Change the color of acid / base indicators React with active metals to release H 2 gas 2HCl + Mg  MgCl 2 + H 2

3. Properties of Acids React with bases to produce salts and H 2 O (neutralization) Conduct electricity - electrolytes

4. Acid Nomenclature Binary Acid - contain H and 1 other element. Prefix “hydro” Root “name of 2nd element” Suffix “ic” Ex. Hydrochloric Acid HCl Hydroiodic Acid - HI

5. Oxyacid H, O and a 3rd element (mostly non-metal) Root - name the oxyion, replace suffix Suffix “ic” for “ate” “ous” for “ite” ions Carbonic Acid H 2 CO 3 Sulfuric Acid H 2 SO 4 2 H + + SO 4 2- sulfate Sulfurous Acid H 2 SO 3 2 H + + SO 3 2- sulfite

6. Common household substances that contain acids and bases. Vinegar is a dilute solution of acetic acid. Drain cleaners contain strong bases such as sodium hydroxide.

7. Common Industrial Acids H 2 SO 4 - sulfuric Acid - most produced chemical in the world. Metallurgy, fertilizer, paper, paint, dyes HNO 3 nitric acid- suffocating odor, stains, protein yellow - explosives, rubber, plastics

8. Common Industrial Acids H 3 PO 4 - phosphoric acid - fertilizer, flavors beverages, cleaning agent HCl - hydrochloric acid - found in stomach, cleaning agent, food processing, dilute forms called “muriatic acid” in stores.

9. Common Industrial Acids Acetic Acid- CH 3 COOH - foul smelling, vinegar - 4% - 6% acetic acid : plastics, food additives.

10. Bases Aqueous solutions taste bitter Change color of acid/ base indicators Dilute aqueous solutions feel slippery

11. Bases React with acids to produce salts and water (neutralization) Conduct electricity - electrolyte

13. Models of Acids and Bases Arrhenius Concept: Acids produce H + in solution, bases produce OH  ion. Brønsted-Lowry: Acids are H + donors, bases are proton acceptors. HCl + H 2 O  Cl  + H 3 O + acid base

14. Arrhenius Definition Acid – produces H + (H 3 O + )in water HCl + H 2 O  H 3 O + + Cl - H 2 SO 4 + H 2 O  2H 3 O + + SO 4 2- Base – produces OH - in water NaOH  Na + + OH - NH 4 OH  NH 4 + + OH -

15. Strong Acid ionizes completely in water - strong electrolyte. Ex. H 2 SO 4, HClO 4, HCl, HNO 3 HBr, HI

16. Weak Acid partially ionizes - weak electrolyte. HCN + H 2 O  H 3 O + + CN – Ex. H 3 PO 4, HF, CH 3 COOH, H 2 CO 3, H 2 S

17. Alkaline Arrhenius bases- increase concentration of OH- in aqueous solution. NaOH  Na + + OH – NH 3 + H 2 O  NH 4 + + OH -

18. Strong Base strong electrolyte - completely dissociate. Ex. NaOH, KOH, LiOH

19. Weak Base Weak electrolyte- partially ionizes; Most of it stays in its molecular form Ex. NH 3 + H 2 O ↔ NH 4 + + OH -

20. Acid/Base Strength Strong Acids HNO 3 HClO 4 HCl HBr HI H 2 SO 4 Strong Bases NaOH KOH LiOH Ba(OH) 2 Ca(OH) 2 (slightly soluble) Sr(OH) 2 (slightly soluble)

21. Bronsted - Lowry Acids and Bases Based on whether a substance is a proton acceptor or donor in non-aqueous solutions. HCl + NH 3  NH 4 + + Cl - HCl donated a proton(H +) to NH 3. Proton donor - acid The NH 3 accepted a proton from HCl proton acceptor- base

22. Bronsted - Lowry Acids and Bases Bronsted-Lowry is a way to study proton transfer!! Acid – H+ donor Base – H+ acceptor For Example: HCl + NH 3  NH 4 + + Cl - H 2 SO 4 + H 2 O  2 H 3 O + + SO 4 2-

23. Conjugate Acids/Bases Conjugate Acid – formed when BL base gains a proton Conjugate Base – formed when BL acid looses a proton HCl + NH 3  NH 4 + + Cl - H 2 SO 4 + 2H 2 O  2 H 3 O + + SO 4 2-

24. Conjugate Acids/Bases HCl + NH 3 <  NH 4 + + Cl - H 2 SO 4 +2H 2 O <  2 H 3 O + + SO 4 2- AcidBaseConj. Acid Conj. Base

25. Conjugates Strength The stronger the acid, the weaker its conjugate base; the stronger the base, the weaker its conjugate acid. Proton transfers favor the production of weaker acids and weaker base. Therefore, CH 3 COOH + Water  H 3 O + + CH 3 COO - (weak acid) (weak base)(stronger acid) (stronger base) Reactants are favored!!

26. Amphoteric can react as either an acid or base HCl + Water  H 3 O + + Cl - proton acceptor(water) Water + NH 3  NH 4 + + OH - Water(proton donor)

27. Monoprotic Acids Ionization – when ions are formed from solute molecules by the action of the solvent. Monoprotic donates 1 hydrogen For example: H 2 O (l) + HCl (s)  H 3 O + (aq) + Cl - (aq) Hydronium ion = H +

28. Polyprotic Acids/Bases Some acids have more than one ionizable hydrogen and are called polyprotic: diprotic (2 H + ), triprotic (3 H + ). For example: 2 H 2 O (l) + H 2 SO 4 (s)  2 H 3 O + (aq) + SO 4 2- (aq) Two moles of hydronium ions

29. Polyprotic Acids Ionization is in several distinct steps: e.g., H 2 CO 3 : carbonic acid H 2 CO 3 + H 2 O  H 3 O + + HCO 3 - HCO 3 - + H 2 O  H 3 O + + CO 3 2- Transfer of 2 nd (or 3 rd ) proton are more difficult than 1 st.

30. Lewis Acids and Bases based on bonding structures - includes acids that do not contain H. Broadest definition of acid/base Includes all BL acids and bases

31. Lewis Acids and Bases Lewis Acid: electron pair acceptor Lewis Base: electron pair donor

32. Lewis Acid an atom, ion or molecule that accepts an electron pair to form a covalent bond. BF 3 + F -  BF 4 (Lewis acid) (Lewis Base)

33. Lewis Base an atom, ion or molecule that donates an electron pair to form a covalent bond.

34. Self-ionization of water a Very weak electrolyte. Autoionization of water: 2 H 2 O (l)  H 3 O + (aq) + OH - (aq) hydronium hydroxide When protons (H+) are produced in water, they bind to the lone pair e- of water to produce H 3 O +

35. Acid/Base Equilibria At 25 o C pure water, has a pH = 7 [H 3 O + ] = [OH - ] = 1 x 10 -7 K w = ionization constant for water K w = 1.0 X 10 -14 @ 25 o C Note: your book presents the autoionization of water on the reaction: H 2 O  H + + OH - ; [H + ] is analogous to [H 3 O + ] !

36. Acids and Bases H 2 O = HOH = H + + OH - AcidsBases HClNaOH HNO 3 KOH HFNH 4 OH

37. The pH Scale pH   log[H + ] pH in water ranges from 0 to 14. K w = 1.00  10  14 = [H + ] [OH  ] pH + pOH = 14.00 As pH rises, pOH falls (sum = 14.00).

38. pH Scale pH – related to the concentration of H + ions in solutions. The more H + ions, the lower the pH.

39. The pH scale and pH values of some common substances.

40. pH Practice 151x10 -14 131x10 -12 111x10 -10 91x10 -8 71x10 -7 51x10 -6 31x10 -4 A1x10 -1 1A21x10 -2 AorB[H 3 O + ]pHAorBpH[H 3 O + ] pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

41. pH Practice B1x10 -15 15B141x10 -14 B1x10 -13 13B121x10 -12 B1x10 -11 11B101x10 -10 B1x10 -9 9B81x10 -8 N1x10 -7 7N7 A1x10 -5 5A61x10 -6 A1x10 -3 3A41x10 -4 A1x10 -1 1A21x10 -2 AorB[H 3 O + ]pHAorBpH[H 3 O + ] pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

42. More pH Practice 4.60x10 -8 5.00x10 -2 7.3x10 -13 3.50x10 -5 2.00x10 -10 A41.00x10 -4 AorBpH[H 3 O + ] pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

43. More pH Practice 4.60x10 -8 5.00x10 -2 7.3x10 -13 3.50x10 -5 2.00x10 -10 A41.00x10 -4 ANBpH[H 3 O + ] pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH 9.7B 4.5A 12.1B 1.3A 7.3B

44. More pH Practice 4.6bananas 7.8eggs 2.0stomach acid 8.5seawater 10.5milk of mag. 3.1apples N4.0x10 -8 7.4blood AorB[H 3 O + ]pHItem pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

45. More pH Practice 4.6bananas 7.8eggs 2.0stomach acid 8.5seawater 10.5milk of mag. 3.1apples N4.0x10 -8 7.4blood AorB[H 3 O + ]pHItem pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH 7.9x10 -4 A 3.2x10 -11 B 3.2x10 -9 B 1.0x10 -2 A 1.6x10 -8 B 2.5x10 -5 A

46. Kw and pOH Practice 4.0x10 -10 2.5x10 -5 4.6bananas 6.25x10 - 7 1.6x10 -8 7.8eggs 1.0x10 -12 1.0x10 -2 2.0stomach acid 3.1x10 -6 3.2x10 -9 8.5seawater 3.1x10 -4 3.2x10 -11 10.5milk of mag. 1.27x10 -11 7.9x10 -4 3.1apples 2.5x10 -7 4.0x10 -8 7.4blood [OH - ][H 3 O + ]pHItem pOH = -log [OH - ]; Kw = [H 3 O + ][OH - ]; Kw = [H 3 O + ][OH - ] = 1 x 10 -14

47. Titrations Titration – experimental technique that provides a sensitive means of determining the chemically equivalent amounts of acid and base. Titration equation:

48. Titrations Equivalence point – point in reaction when equal # moles of acid and base have reacted. Neutralization reaction equation: HCl + NaOH  NaCl + H 2 O 1 mol acidbasesaltwater

49. Titrations We can now look at titrations more quantitatively. Possible combinations we will consider: strong acid (SA) - strong base (SB) weak acid (WA) - strong base (WB) strong acid (SA) - weak base (WB) Consider each case before/at/after equivalence point (EP).

50. Titrations A Titration Curve shows pH at various points in a titration experiment(before, at, and after equivalence point (EP)). Can generate by: experimentally measuring pH during a titration experiment calculating pH at various points for an acid-base reaction

51. Titrations The Titration curve of acids with bases (or visa versa) has four different environments to consider: 1.Before rxn (before titration begins) 2.Before the EP (at the very beginning of the titration) 3.At the EP (moles of acids = moles of base) 4.After the EP

52. SA/SB Titrations KOH (aq) + HBr (aq)  KBr (aq) + H 2 O (l) Net reaction: OH - (aq) + H + (aq)  H 2 O @ Equivalence Point (EP): What compounds are present at EP? H 2 O, K +, and Br - K + and Br - are conjugates of strong base and acid; too weak to react with water! @ EP, [OH - ] = [H 3 O + ]; pH=7

53. SA/SB Titrations The Titration Curve for a SA/SB will look like this. But: Acid/Base “neutralization” reactions do not always result in “neutral” solutions!

54. SA/SB Titrations Before & After Equivalence Point (EP) In this case, we need to consider what species are present and at what concentrations. At each stage – we need the [H 3 O + ] or [OH - ] concentration! Example: KOH + HBr Have 20.0 ml 2.0 M HBr; titrate with 1.0 M KOH

55. SA/SB Titrations 20.0 ml 2.0 M HBr; titrate 1.0 M KOH Before rxn: 2.0 M HBr = 2.0 M H 3 O + pH = -0.30 ~ 0 0.0200 L (2.0 M) = 0.040 mol HBr present M A V A = M B V B (2.0 M)(0.02 L) = (1.0 M)(vol KOH) EP will occur at 0.04 L or 40 ml KOH

56. SA/SB Titrations 20.0 ml 2.0 M HBr; titrate 1.0 M KOH Before Equivalence Point (30.0 ml KOH): 30.0 ml KOH = 0.0300 L = 0.0300 mol KOH added 0.0300 mol HBr have reacted 0.040-0.0300 = 0.010 mol HBr in 50.0 ml  0.20 M H 3 O + pH = +0.70

57. SA/SB Titrations 20.0 ml 2.0 M HBr; titrate 1.0 M KOH @ Equivalence Point (40.0 ml KOH): 40.0 ml KOH = 0.0400 L = 0.0400 mol KOH added 0.0400 mol HBr have reacted 0.040-0.0400 = 0.000 mol HBr in 60.0 ml  [OH - ] = [H 3 O + ]; pH = 7

58. SA/SB Titrations 20.0 ml 2.0 M HBr; titrate 1.0 M KOH After Equivalence Point (50.0 ml KOH): 50.0 ml KOH = 0.0500 L = 0.0500 mol KOH added all HBr reacted; 0.010 mol KOH leftover! 0.010 mol KOH in 70.0 ml  [OH - ] = 0.1428 M OH - pOH = 0.84; pH = 13.15

59. WA/SB Titrations F - + H 2 O  HF + OH - This equilibrium will make the solution basic. The EP of a WA/SB reaction is always basic.

60. SA/WB Titrations NH 4 + + H 2 O  H 3 O + + NH 3 This equilibrium will make the solution acidic. The EP of a SA/WB reaction is always acidic.

61. Titration Summary Strong Acid/ Strong Base

62. Titration Summary Weak Acid/ Strong Base

63. Titration Summary Strong Acid/ Weak Base

64. Titration Summary

65. Polyprotic Acid Titrations When polyprotic acids are titrated with strong bases, there are multiple equivalence points. The titration curve of a polyprotic acid shows an equivalence point for the each acid H + :

66. Acid-Base Properties of Salts