1. Chapter 17 Acids and Bases

2. Acids, Bases, and Matter Classification of Matter

3. Arrhenius Definition of Acids and Bases Acids Acidus – Latin for sour Substances that, when added to water, produce hydrogen ions, H+ (protons), or hydronium ions, H 3 O +.

4. Bases Alkali – Arabic for ashes. Product of burning plants is potash (KOH) which feels soapy and tastes bitter. Substance that when added to water produce hydroxide ions, OH -.

5. Salts Substance formed in addition to water when an acid reacts with a base. The cation is from the base and the anion is form the acid.

6. Limitations of Arrhenius Definition Arrhenius's concept is limited to aqueous solutions because it refers to ions derived from water.

7. The Hydronium Ion and Water Autoionization There is an equilibrium between these two ions in water or any aqueous solution: H 2 O + H 2 O H 3 O + (aq) + OH - (aq)

8. . The Bronsted Concept of Acids and Bases Classifies an acid as a proton donor and a base as a proton acceptor. Another way to look at the model is an acid is a substance from which a proton can be removed and a base is a substance that can remove a proton from an acid.

9. Acids Recognizing acids is fairly simple - since an acid has to be able to "donate" a proton it must contain an ionizable hydrogen. Generally speaking this means that an acid is something that begins with a hydrogen. It can be hydrogen attached to a polyatomic ion or hydrogen attached to a nonmetal.

10. Acids Continued Acids that are capable of donating one proton are called monoprotic acids. Some acids are capable of donating more than one proton and are called polyprotic acids.

11. Bases Bases are not so simple- in most cases bases will be ionic compounds containing the OH- ion. When the compound is dissolved in water the OH- ion dissociates from the positive ion and is then able to "accept" the proton. In other cases regular anions act as bases. The only bases that don't fit into these categories are ammonia and derivatives of ammonia. The ammonia molecule - for reasons which aren't important to us - is able to "accept" a proton and become the ammonium ion, NH 4 +.

12. Bases Continued Bases can also be monoprotic or polyprotic depending on how many protons they accept.

13. Amphiprotic substances These are molecules or ions that can behave as a Bronsted acid or base. One of the best examples of this is water.

14. Conjugate Acid-Base Pairs A pair of compounds or ions that differ be the presence of one H + unit is called a conjugate acid-base pair. When you have the proton you are the acid. If you do not you are the conjugate base.

15. Conjugate Acid-Base Pairs Every acid-base reaction involving H + transfer has two conjugate acid-base pairs.

16. Relative Strengths of Acids and Bases The strength of an acid in solution has to do with the relative ability of an acid to donate protons, where an acid is considered strong only if it is such a good proton donor that each and every acid molecule will give up at least one hydrogen ion to a water molecule (100% acid dissociations). (The stronger the acid the weaker it holds the proton.)

17. In the Bronsted model, an acid donates a proton and produces a conjugate base. this model also informs us that, in general, the stronger the acid, the weaker it conjugate base.

18. Strong Acids Completely ionize in water to form H 3 O + ions: HX(aq) + H 2 O(l) ----------> H 3 O + (aq) + X - (aq) Include: HCl, HBr, HI, HNO 3, HClO 4, H 2 SO 4

19. Strong Bases Completely ionize to form OH - ions in solution. Hydroxides of the Group I metals: NaOH(s) ----------> Na + (aq) + OH - (aq) Hydroxides of the heavier Group II metals M(OH) 2 (s) ----------> M 2+ (aq) + 2OH - (aq) M = Ca, Sr, Ba

20. Weak Acids Partially dissociate in water to form H 3 O + ions: HX(aq) + H 2 O(l) H 3 O + (aq) + X - (aq) Generally, concentration of HX molecules >> concentration of H 3 O + ions. Concentrations are governed by equilibrium constants - Ka.

21. General Equation HB(aq) + H 2 O(l) H 3 O + (aq) + B - (aq) Ka = [H 3 O + ] [B - ] / [HB] The smaller the ionization constant, the weaker the acid. See table 17.4 pg 808

22. Types of Weak Acids Molecular Weak Acids (acid-base indicators are molecular weak acid) HIn(aq) + H 2 O(l) -----> H 3 O + (aq) + In - (aq) blue Yellow

23. Types of Weak Acids Anions containing an ionizable H atom HSO 4 - (aq), HCO 3 -, H 2 PO 4 -, etc

24. Types of Weak Acids All cations except those of Group I metals, Ca 2+, Sr 2+, Ba 2+ Ammonium, NH 4 + Zn(H 2 O) 4 2+ + H 2 O(l) H 3 O + (aq) + Zn(H 2 O) 3 OH +

25. Weak Bases Concentrations are governed by equilibrium constants - Kb. General Equation B - (aq) + H 2 O(l) HB(aq) + OH - (aq) Kb = [HB] [OH - ] / [B - ] The smaller the ionization constant, the weaker the base.

26. Types of Weak Bases Molecular: NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq)

27. Types of Weak Bases Anions derived from weak acids (conjugate bases): F - (aq) + H 2 O(l) HF(aq) + OH - (aq)

28. Predict the Predominate Direction of Acid-Base Reactions The following chart can be used to predict whether the equilibrium in an acid-base reaction lies predominately to the left or the right. Examples. 1. CH 3 COOH and NaCN 2. NH 4 Cl and Na 2 CO 3

30. The Water Ionization Constant, Kw Kw = [H 3 O + ] [OH - ] = 1.00 x 10 -14 at 25 C

31. Water, Acids and Bases Pure Water[H 3 O + ]= [OH - ] = 1.00 x 10 -7 Acidic [H 3 O + ] > [OH - ] [H 3 O + ] > 1.00 x10 -7 Basic [H 3 O + ] < [OH - ] [H 3 O + ] < 1.00 x10 -7

32. Autoionization of Water All acids/bases dissolved in water must obey equation for the ionization of water. –They either add H 3 O + or OH  to water. Most of the acids in this chapter will be stronger than water and add significantly to the hydronium ion concentration.

33. Examples The hydronium ion concentration of an acidic solution was 1.00x10  5 M. What was the [OH  ]? What is the hydronium ion concentration if the hydroxide concentration was 2.50x10  3 M?

34. The pH Scale pH = - log 10 [H + ] = - log 10 [H 3 O + ] Relation to Aqueous Solution Neutral: pH = 7.00 Acidic:pH < 7.00 Basic:pH > 7.00

35. Methods of Measuring pH pH paper is used that has compounds in it which are change to different colors for different pH ranges.

36. Methods of Measuring pH An colored indicator can be placed in the solution and its color correlated with pH. HIn(aq) + H 2 O(l)  H 3 O + (aq) + In  (aq). E.g. phenolphthalein is colorless in acid form but pink in basic form. –The pH at which they change color depends on their equilibrium constant.

37. Methods of Measuring pH More accurate and precise measurements are made with a pH meter. A combination of voltmeter and electrodes

38. The pOH pOH= - log 10 [OH - ] Given: pKw = - log 10 1.00 x 10 -14 = 14 pKw = 14 = pOH + pH

39. Examples Determine the pH of a solution in which [H 3 O + ] = 5.40x10  6 M Determine the pH of a solution in which the [OH  ] = 3.33x10  3 M

40. Examples Determine the pOH of a solution in which the [OH  ] = 3.33x10  3 M Determine the [H 3 O + ] if the pH of the solution is 7.35.

41. The Ionization Constants for an Acid and Its Conjugate Base Ka x Kb = [H 3 O + ] [B - ] / [HB] x [HB] [OH - ] / [B - ] = [H 3 O + ] [OH - ] = 1.00x 10 -14 = Kw The strength of the base is inversely related to that of its conjugate weak acid.

42. Equilibria Involving Weak Acids Calculating Ka or Kb from Initial Concentrations and Measured pH Calculation of Ka (in water) [H 3 O + ] = [B - ] = antilog –pH [HB] = original concentration - [H 3 O + ] Percent Ionization Percent Ionization HB = [H 3 O + ] / original concentration HB

43. Weak Acids and the ICE Table

44. Equilibria Involving Weak Bases Same procedure as with Ka (H + and OH - exchange)

45. Weak Bases and the ICE Table

46. Calculating Equilibrium Concentrations and pH from Initial Conc. and Ka or Kb Dissolving a weak acid, HB, in water [H 3 O + ] = [B - ] [HB] = original concentration - [H 3 O + ]

47. Making it easier In general, Ka is seldom known to better than +/- 5%, Hence, in the expression Ka = [H 3 O + ] [B - ] / [HB] = x 2 /(a-x) if x/a < 0.05 we take a - x = a

48. Calculation of [OH-] in Solution of Weak Base Same procedure as the calculation of [H 3 O + ] in solution of weak acid.

49. Acid-Base Properties of Salts Salt A salt is an ionic compound that COULD have been formed by the reaction of an acid with a base; a salt's positive ions come from the base, and its negative ions come from the acid.

50. Hydrolysis reaction A hydrolysis reaction is said to have occurred when a salt dissolves in water and leads to changes in the hydronium or hydroxide ion concentrations of the water. To determine whether a salt is acidic or basis you must consider the effect of the cation and the anion separately and then combine these effects to give the overall result for the salt.

51. Cations Neutral derived from strong bases Li+, Na+, K+, Ca2+, Sr2+, Ba2+ Acidic all other cations, including those of transition metals

52. Anions Neutral derived from strong acids HSO 4 -, H 2 PO 4 -, etc Basic all other anions. In general, an anion derived from a weak acid is expected to be basic

53. Overall Result

54. Salts of Weak Acids and Weak Bases Ka > Kbacidic salt Ka = Kbneutral salt Ka < Kbbasic salt

55. Polyprotic Acids and Bases Ionization Steps

56. Polyprotic Acids and Bases Calculating the pH of the Solution

57. Molecular Structure, Bonding, and Acid-Base Behavior One of the most interesting aspects of chemistry is the correlation between molecular structure, bonding, and observed properties. Acid Dissociation Steps 1. H-A bond breaking 2. Loss of an electron by H to form H+ 3. Gain of an electron by A to form A-.

58. Because the second step is common to all acid dissociations the relative strength of an acid must be determined from steps one and three. Step one is a measure of bond strength, the weaker the bond the easier it is for the H and A to break apart. Step three is a measure of electron affinity, the higher the electron affinity, the stronger the attraction of the atom A for the electron.

59. Oxyacids and the Inductive effect The greater the number of oxygen atoms attached to the central atom in the acid, the stronger the acid. This increase in acidity is due to the inductive effect, the attraction of electrons from adjacent bonds by the more electronegative atom. As more and more O atoms are attached, the inductive effect is stronger, the O-H bond id more strongly polarized and the bond is more readily broken.

60. Why All Hydrogens Do Not Ionize The cleavage of a bond to release a hydrogen ion can only occur if the attached atom can accommodate the negative charge.

61. The Lewis Concept of Acids and Bases Definitions Acid: electron-pair acceptor (must have < octet or be positively charged) Base: electron-pair donor (must have nonbonding electrons available for coordinate covalent bonding)

62. Lewis Acids Cationic Acids All metal cations are potential Lewis acids. Molecular Acids Oxides of nonmetals

63. Lewis Bases Must have a pair of electrons not used in bonding

64. Comparison of Theories Lewis definition most general, then Brønsted-Lowery and finally Arhennius: