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 + in water HCl  H + + Cl - H 2 SO 4  2H + + 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 http://www.mhhe.com/physsci/chemistry/chang7/esp/folder_structure/ac/m2/s1/acm2s1_1.htm

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 Answer questions on strong or weak (under the picture) ! http://www.elmhurst.edu/~chm/vchembook/185strength.html

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.

41. HIn --> H+ + In - For phenolphthalein: pH 0 to 8.2 = colorless; then pink, then pH 10 = red Add H+ then shift to left Add OH- then shift to right

42. 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

43. 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

44. 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

45. More pH Practice B7.34.60x10 -8 A1.35.00x10 -2 B12.17.3x10 -13 A4.53.50x10 -5 B9.72.00x10 -10 A41.00x10 -4 ANBpH[H 3 O + ] pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

46. 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

47. More pH Practice A2.5x10 -5 4.6bananas B1.6x10 -8 7.8eggs A1.0x10 -2 2.0stomach acid B3.2x10 -9 8.5seawater B3.2x10 -11 10.5milk of mag. A7.9x10 -4 3.1apples N4.0x10 -8 7.4blood AorB[H 3 O + ]pHItem pH = -log [H 3 O + ];[H 3 O + ] = 10 -pH

48. 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

49. Titrations Titration – experimental technique that provides a sensitive means of determining the chemically equivalent amounts of acid and base. Titration equation: How to titrate !! http://www.wwnorton.com/chemistry/tutorials/ch16.htm Vacid x (Molarity Acid x # Equivalent acid) = Vbase x (Molarity base x # Eq base)

50. 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

51. 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).

52. 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

53. 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

54. 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

55. 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!

56. 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

57. 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

58. 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

59. 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

60. 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

61. 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.

62. 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.

63. Titration Summary Strong Acid/ Strong Base

64. Titration Summary Weak Acid/ Strong Base

65. Titration Summary Strong Acid/ Weak Base

66. Titration Summary

67. http://www.physchem.co.za/Acids/Titrations.htm Click on Titration curve Then, you pick the correct Curve in the question

68. 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 + :

69. Acid-Base Properties of Salts