2. Arrhenius Definition Acids produce hydrogen ions in aqueous solution. Acids produce hydrogen ions in aqueous solution.  H 2 SO 4, HCl, HC 2 H 3 O 2 Bases produce hydroxide ions when dissolved in water. Bases produce hydroxide ions when dissolved in water.  NaOH, KOH This definition is limited to aqueous solutions. This definition is limited to aqueous solutions. Only one kind of base (a producer of OH - ). Only one kind of base (a producer of OH - ).  NH 3 (ammonia) is not a base using this definition.

3. Bronsted-Lowry Definitions And acid is an proton (H + ) donor and a base is a proton acceptor. And acid is an proton (H + ) donor and a base is a proton acceptor. Acids and bases always come in pairs. Acids and bases always come in pairs. HCl is an acid.. HCl is an acid.. When it dissolves in water it gives its proton to water. When it dissolves in water it gives its proton to water. HCl (g) + H 2 O (l) H 3 O + (aq) + Cl - (aq) HCl (g) + H 2 O (l) H 3 O + (aq) + Cl - (aq) Water is the base in the acid dissociation Water is the base in the acid dissociation  When water accepts proton it produces a hydronium ion (H 3 O + )

4. Pairs Acids are always paired with conjugate bases Acids are always paired with conjugate bases General equation General equation HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) Acid + Base Conjugate acid (CA) + Conjugate base (CB) Acid + Base Conjugate acid (CA) + Conjugate base (CB)  This is an equilibrium. There is competition for H + between H 2 O and A - There is competition for H + between H 2 O and A -  The stronger base controls side that is favored. If H 2 O is a stronger base than A - it takes the H + If H 2 O is a stronger base than A - it takes the H +  Equilibrium moves to right.

5. Acid dissociation constant K a The equilibrium constant for the general equation. The equilibrium constant for the general equation. HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) HA (aq) + H 2 O (l) H 3 O + (aq) + A - (aq) K a = [H 3 O + ][A - ] [HA] K a = [H 3 O + ][A - ] [HA] H 3 O + is often written H + ignoring the water in equation (it is implied). H 3 O + is often written H + ignoring the water in equation (it is implied).

6. Acid dissociation constant K a HA (aq) H + (aq) + A - (aq) HA (aq) H + (aq) + A - (aq) K a = [H + ][A - ] [HA] K a = [H + ][A - ] [HA] We can write the expression for any acid. We can write the expression for any acid. Strong acids dissociate completely. Strong acids dissociate completely.  Equilibrium lies far to product side. The conjugate base of a strong acid must be very weak. The conjugate base of a strong acid must be very weak.

7. Back to Pairs Strong acids Strong acids K a is large K a is large [H + ] is equal to [HA] dissolved [H + ] is equal to [HA] dissolved A - is a weaker base than water A - is a weaker base than water Weak acids K a is small [H + ] << [HA] A - is a stronger base than water

8. Types of Acids Polyprotic Acids- more than 1 acidic hydrogen (diprotic, triprotic). Polyprotic Acids- more than 1 acidic hydrogen (diprotic, triprotic). Oxyacids - Proton is attached to the oxygen of an ion. Oxyacids - Proton is attached to the oxygen of an ion. Organic acids contain the Carboxyl group -COOH with the H attached to O Organic acids contain the Carboxyl group -COOH with the H attached to O  Organic acids are generally very weak.

9. Amphoteric Compounds Compounds that behave as both an acid and a base. Compounds that behave as both an acid and a base. Water autoionizes Water autoionizes 2H 2 O (l) H 3 O + (aq) + OH - (aq) 2H 2 O (l) H 3 O + (aq) + OH - (aq) K w = [H 3 O + ][OH - ]=[H + ][OH - ] K w = [H 3 O + ][OH - ]=[H + ][OH - ] At 25ºC K w = 1.0 x10 -14 At 25ºC K w = 1.0 x10 -14 In EVERY aqueous solution @ 25 o C. In EVERY aqueous solution @ 25 o C.  Neutral solution [H + ] = [OH - ]= 1.0 x10 -7  Acidic solution [H + ] > [OH - ]  Basic solution [H + ] < [OH - ]

10. pH, pOH and pK a pH= -log[H + ] pH= -log[H + ] Used because [H + ] is usually very small Used because [H + ] is usually very small As pH decreases, [H + ] increases exponentially As pH decreases, [H + ] increases exponentially [H + ] = 1.0 x 10 -8 pH= 8.00 [H + ] = 1.0 x 10 -8 pH= 8.00 pOH= -log[OH - ] pOH= -log[OH - ] pK a = -log K a pK a = -log K a

11. Relationships K W = [H + ][OH - ] K W = [H + ][OH - ] -log K w = -log([H + ][OH - ]) -log K w = -log([H + ][OH - ]) -log K w = -log[H + ]+ -log[OH - ] -log K w = -log[H + ]+ -log[OH - ] pK w = pH + pOH pK w = pH + pOH K w = 1.0 x10 -14 K w = 1.0 x10 -14 pK W = 14 pK W = 14 14.00 = pH + pOH 14.00 = pH + pOH [H + ], [OH - ], pH and pOH [H + ], [OH - ], pH and pOH  Given any one of these we can find the other three.

12. BasicAcidicNeutral 10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14 [H + ] 013579111314 pH Basic 10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14 [OH - ] 013579111314 pOH

13. Calculating pH of Solutions Always write down the major ions in solution Always write down the major ions in solution  This first step is the most important Remember these are equilibria Remember these are equilibria  You need to consider the direction Remember the chemistry Remember the chemistry Don’t try to memorize individual cases Don’t try to memorize individual cases  Technique is approximately the same

14. Strong Acids HBr, HI, HCl, HNO 3, H 2 SO 4, HClO 4 HBr, HI, HCl, HNO 3, H 2 SO 4, HClO 4 ALWAYS WRITE THE MAJOR SPECIES ALWAYS WRITE THE MAJOR SPECIES These are always completely dissociated These are always completely dissociated  No equilibrium [H + ] = [HA] dissolved [H + ] = [HA] dissolved [OH - ] is going to be small because of equilibrium [OH - ] is going to be small because of equilibrium  10 -14 = [H + ][OH - ] If [HA]< 10 -7 water contributes H + If [HA]< 10 -7 water contributes H +

15. Weak Acids K a will be small. K a will be small. ALWAYS WRITE THE MAJOR SPECIES. ALWAYS WRITE THE MAJOR SPECIES. Approach as an equilibrium problem from the start. Approach as an equilibrium problem from the start. Determine whether most of the H + will come from the acid or the water. Determine whether most of the H + will come from the acid or the water.  Compare K a or K w Solve the problem like a normal equilibrium Solve the problem like a normal equilibrium

16. Example Calculate the pH, pOH and [OH - ] of 2.0 M acetic acid HC 2 H 3 O 2 with a K a 1.8 x10 -5 Calculate the pH, pOH and [OH - ] of 2.0 M acetic acid HC 2 H 3 O 2 with a K a 1.8 x10 -5

17. Example Calculate the pH, pOH and [OH - ] of 0.15 M iodic acid HIO 3 with a K a 1.7 x10 -1 Calculate the pH, pOH and [OH - ] of 0.15 M iodic acid HIO 3 with a K a 1.7 x10 -1

18. Polyprotic acids Always dissociate stepwise. Always dissociate stepwise. The first H + comes of much easier than the second. The first H + comes of much easier than the second. Ka for the first step is bigger than Ka for the second. Ka for the first step is bigger than Ka for the second. Denoted Ka 1, Ka 2, Ka 3 Denoted Ka 1, Ka 2, Ka 3

19. Example Calculate the pH, pOH and [OH - ] of 0.750 M phosphoric acid H 3 PO 4 with a Calculate the pH, pOH and [OH - ] of 0.750 M phosphoric acid H 3 PO 4 with a  K a1 = 7.50 x 10 -3  K a2 = 6.20 x 10 -8  K a3 = 4.20 x 10 -13

20. Calculate the Concentration Of all the ions in a solution of 1.00 M Arsenic acid H 3 AsO 4 Of all the ions in a solution of 1.00 M Arsenic acid H 3 AsO 4 n Ka 1 = 5.0 x 10 -3 n Ka 2 = 8.0 x 10 -8 n Ka 3 = 6.0 x 10 -10

21. A mixture of Weak Acids The process is the same. The process is the same. Determine the major species. Determine the major species. The stronger acid will always predominate. The stronger acid will always predominate.  Ignore the weaker acid(s)  Use the bigger K a if concentrations are comparable Calculate the pH of a mixture 1.20 M HF (Ka = 6.8 x 10 -4 ) and 3.4 M HOC 6 H 5 (Ka = 1.3 x 10 -10 ) Calculate the pH of a mixture 1.20 M HF (Ka = 6.8 x 10 -4 ) and 3.4 M HOC 6 H 5 (Ka = 1.3 x 10 -10 )

22. Percent dissociation = amount dissociated x 100 initial concentration = amount dissociated x 100 initial concentration For a weak acid percent dissociation increases as acid becomes more dilute. For a weak acid percent dissociation increases as acid becomes more dilute. Calculate the % dissociation of 1.00 M Acetic acid (Ka = 1.8 x 10 -5 ) Calculate the % dissociation of 1.00 M Acetic acid (Ka = 1.8 x 10 -5 ) Calculate the % dissociation of 0.0100 M Acetic acid (Ka = 1.8 x 10 -5 ) Calculate the % dissociation of 0.0100 M Acetic acid (Ka = 1.8 x 10 -5 ) As [HA] o decreases [H + ] decreases As [HA] o decreases [H + ] decreases  but % dissociation increases. Relate this to Le Chatelier Relate this to Le Chatelier

23. The other way What is the Ka of a weak acid that is 8.1 % dissociated as 0.100 M solution? What is the Ka of a weak acid that is 8.1 % dissociated as 0.100 M solution?

24. Bases The OH - is a strong base. The OH - is a strong base. Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved. Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved. The hydroxides of alkaline earths Ca(OH) 2 etc. are strong dibasic bases, but they don’t dissolve well in water. The hydroxides of alkaline earths Ca(OH) 2 etc. are strong dibasic bases, but they don’t dissolve well in water.  The hydroxides of alkaline earths are used as antacids because [OH - ] can’t build up.

25. Bases without OH - Bases are proton acceptors. Bases are proton acceptors. NH 3 + H 2 O NH 4 + + OH - NH 3 + H 2 O NH 4 + + OH - It is the lone pair on nitrogen that accepts the proton. It is the lone pair on nitrogen that accepts the proton. Many weak bases contain nitrogen Many weak bases contain nitrogen B (aq) + H 2 O (l) BH + (aq) + OH - (aq) B (aq) + H 2 O (l) BH + (aq) + OH - (aq) K b = [BH + ][OH - ] [B] K b = [BH + ][OH - ] [B]

26. Strength of Bases N:N: Hydroxides are strong. Hydroxides are strong. Others are weak. Others are weak. Smaller K b weaker base. Smaller K b weaker base. Calculate the pH of a solution of 4.0 M pyridine (K b = 1.7 x 10 -9 ) Calculate the pH of a solution of 4.0 M pyridine (K b = 1.7 x 10 -9 )

27. Polyprotic acid H 2 CO 3 H + + HCO 3 - Ka 1 = 4.3 x 10 -7 H 2 CO 3 H + + HCO 3 - Ka 1 = 4.3 x 10 -7 HCO 3 - H + + CO 3 -2 Ka 2 = 4.3 x 10 -10 HCO 3 - H + + CO 3 -2 Ka 2 = 4.3 x 10 -10 Base in first step is acid in second. Base in first step is acid in second. In calculations we can normally ignore the second dissociation. In calculations we can normally ignore the second dissociation.

28. Sulfuric acid is special In first step it is a strong acid. In first step it is a strong acid. Ka 2 = 1.2 x 10 -2 Ka 2 = 1.2 x 10 -2 Calculate the concentrations in a 2.0 M solution of H 2 SO 4 Calculate the concentrations in a 2.0 M solution of H 2 SO 4 Calculate the concentrations in a 2.0 x 10 -3 M solution of H 2 SO 4 Calculate the concentrations in a 2.0 x 10 -3 M solution of H 2 SO 4

29. Salts as acids an bases Salts are ionic compounds. Salts are ionic compounds. Salts of the cation of strong bases and the anion of strong acids are neutral. Salts of the cation of strong bases and the anion of strong acids are neutral. for example NaCl, KNO 3 for example NaCl, KNO 3 There is no equilibrium for strong acids and bases. There is no equilibrium for strong acids and bases. We ignore the reverse reaction. We ignore the reverse reaction.

30. Basic Salts If the anion of a salt is the conjugate base of a weak acid - basic solution. If the anion of a salt is the conjugate base of a weak acid - basic solution. In an aqueous solution of NaF In an aqueous solution of NaF The major species are Na +, F -, and H 2 O The major species are Na +, F -, and H 2 O F - + H 2 O HF + OH - F - + H 2 O HF + OH - K b =[HF][OH - ] [F - ] K b =[HF][OH - ] [F - ] but Ka = [H + ][F - ] [HF] but Ka = [H + ][F - ] [HF]

31. Basic Salts K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF]

32. Basic Salts K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF]

33. Basic Salts K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF]

34. Basic Salts K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] K a x K b =[OH - ] [H + ] K a x K b =[OH - ] [H + ]

35. Basic Salts K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] K a x K b = [HF][OH - ]x [H + ][F - ] [F - ] [HF] K a x K b =[OH - ] [H + ] K a x K b =[OH - ] [H + ] K a x K b = K W K a x K b = K W

36. K a tells us K b The anion of a weak acid is a weak base. The anion of a weak acid is a weak base. Calculate the pH of a solution of 1.00 M NaCN. K a of HCN is 6.2 x 10 -10 Calculate the pH of a solution of 1.00 M NaCN. K a of HCN is 6.2 x 10 -10 The CN - ion competes with OH - for the H + The CN - ion competes with OH - for the H +

37. Acidic salts A salt with the cation of a weak base and the anion of a strong acid will be basic. A salt with the cation of a weak base and the anion of a strong acid will be basic. The same development as bases leads to The same development as bases leads to K a x K b = K W K a x K b = K W Calculate the pH of a solution of 0.40 M NH 4 Cl (the K b of NH 3 1.8 x 10 -5 ). Calculate the pH of a solution of 0.40 M NH 4 Cl (the K b of NH 3 1.8 x 10 -5 ). Other acidic salts are those of highly charged metal ions. Other acidic salts are those of highly charged metal ions.

38. Anion of weak acid, cation of weak base K a > K b acidic K a > K b acidic K a < K b basic K a < K b basic K a = K b Neutral K a = K b Neutral

39. Structure and Acid base Properties Any molecule with an H in it is a potential acid. Any molecule with an H in it is a potential acid. The stronger the X-H bond the less acidic (compare bond dissociation energies). The stronger the X-H bond the less acidic (compare bond dissociation energies). The more polar the X-H bond the stronger the acid (use electronegativities). The more polar the X-H bond the stronger the acid (use electronegativities). The more polar H-O-X bond -stronger acid. The more polar H-O-X bond -stronger acid.

40. Strength of oxyacids The more oxygen hooked to the central atom, the more acidic the hydrogen. The more oxygen hooked to the central atom, the more acidic the hydrogen. HClO 4 > HClO 3 > HClO 2 > HClO HClO 4 > HClO 3 > HClO 2 > HClO Remember that the H is attached to an oxygen atom. Remember that the H is attached to an oxygen atom. The oxygens are electronegative The oxygens are electronegative

41. Strength of oxyacids Electron Density ClOH

42. Strength of oxyacids Electron Density ClOHO

43. Strength of oxyacids ClOH O O Electron Density

44. Strength of oxyacids ClOH O O O Electron Density

45. Hydrated metals Highly charged metal ions pull the electrons of surrounding water molecules toward them. Highly charged metal ions pull the electrons of surrounding water molecules toward them. Make it easier for H + to come off. Make it easier for H + to come off. Al +3 O H H

46. Acid-Base Properties of Oxides Non-metal oxides dissolved in water can make acids. Non-metal oxides dissolved in water can make acids. SO 3 (g) + H 2 O(l) H 2 SO 4 (aq) SO 3 (g) + H 2 O(l) H 2 SO 4 (aq) Ionic oxides dissolve in water to produce bases. Ionic oxides dissolve in water to produce bases. CaO(s) + H 2 O(l) Ca(OH) 2 (aq) CaO(s) + H 2 O(l) Ca(OH) 2 (aq)

47. Lewis Acids and Bases Most general definition. Most general definition. Acids are electron pair acceptors. Acids are electron pair acceptors. Bases are electron pair donors. Bases are electron pair donors. BF F F :N:N H H H

48. Lewis Acids and Bases Boron triflouride wants more electrons. Boron triflouride wants more electrons. BF F F :N:N H H H

49. Lewis Acids and Bases Boron triflouride wants more electrons. Boron triflouride wants more electrons. BF 3 is Lewis base NH 3 is a Lewis Acid. BF 3 is Lewis base NH 3 is a Lewis Acid. B F F F N H H H

50. Lewis Acids and Bases Al +3 ( ) H H O Al ( ) 6 H H O + 6 +3