1. Topic C – Part I: Acid Base Equilibria and Ksp

2. Arrhenius Definition Acids produce hydrogen ions (H+) in aqueous solution. Bases produce hydroxide ions (OH-) when dissolved in water. Only one kind of base. NH 3 ammonia could not be an Arrhenius base.

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

4. Conjugate Pairs  General equation  HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq)  Acid + Base Conj. acid + Conj. base  This is an equilibrium.  Competition for H + between H 2 O and A -  The stronger base controls direction.  If H 2 O is a stronger base it takes the H +  Equilibrium moves to right.

5. Acid dissociation constant K a The equilibrium constant for the general equation. HA(aq) + H 2 O(l) H 3 O + (aq) + A - (aq) K a = [H 3 O + ][A - ] [HA] 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) K a = [H + ][A - ] [HA] We can write the expression for any acid. Strong acids dissociate completely. Equilibrium far to right. Conjugate base must be weak.

7. Sample Ex. 14.1  Acid Dissociation (ionization) Reactions  Write the simple dissociation reaction (omitting water) for each of the following acids.  Hydrochloric Acid (HCl)  Acetic Acid (HC 2 H 3 O 2 )  The Ammonium Ion (NH 4 + )  The anilinium ion (C 6 H 4 NH 3 + )  The hydrated aluminum (III) ion [Al(H 2 O) 6 ] 3+

8. Describing Acid Strength  Strong acids  K a is large  [H + ] is equal to [HA]  A - is a weaker base than water  Weak acids  K a is small  [H + ] <<< [HA]  A - is a stronger base than water

9. Sample Ex. 14.2 Using table 14.2 (p. 628) arrange the following species according to their strength as bases: H 2 O, F -, Cl -, NO 2 and CN -

10. The concentration of H 3 O+ and OH- are dependent upon two, separate factors 1. The strength of the acid or base (the degree of ionization) 2.The amount of water present (the concentration of the solution). It is possible to have a dilute, (but) strong acid or to have a concentrated, (but) weak acid They might also have the same hydronium ion concentration and pH

11. Examples of Acids and Bases

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

13. Amphoteric  Behave as both an acid and a base.  Water autoionizes  2H 2 O(l) H 3 O + (aq) + OH - (aq)  K W = [H 3 O + ][OH - ]=[H + ][OH - ]  At 25ºC K W = 1.0 x10 -14  In EVERY aqueous solution.  Neutral solution [H + ] = [OH - ]= 1.0 x10 -7  Acidic solution [H + ] > [OH - ]  Basic solution [H + ] < [OH - ]

14. Sample Ex. 14.3 Calculate [H+] or [OH-] as required for each of the following solutions at 25 o C, and state whether the solution is neutral, acidic, or basic. 1.0 x 10 -5 M OH - 1.0 x 10 -7 M OH - 10.0 M H +

15.  pH= -log[H + ]  Used because [H + ] is usually very small  As pH decreases, [H + ] increases exponentially  Sig figs only the digits after the decimal place of a pH are significant  [H + ] = 1.0 x 10 -8 pH= 8.00 2 sig figs  pOH= -log[OH - ]  pKa = -log K pH

16. Sample Ex. 14.5 Calculating pH and pOH Calculate pH and pOH for each of the following solutions at 25 o C. a. 1.0 x 10 -3 M OH - b. 1.0 M H +

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

18. Sample Ex. 14.6 The pH of a sample of human blood was measured to be 7.41 at 25 o C. Calculate pOH, [H + ], and [OH - ] for the sample.

19. 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

20. Calculating pH of Solutions Always write down the major ions in solution. Remember these are equilibria. Remember the chemistry. Don’t try to memorize there is no one way to do this.

21. Strong Acids HBr, HI, HCl, HNO 3, H 2 SO 4, HClO 4 ALWAYS WRITE THE MAJOR SPECIES Completely dissociated [H + ] = [HA] [OH - ] is going to be small because of equilibrium 10 -14 = [H + ][OH - ] If [HA]< 10 -7 water contributes H +

22. Sample Ex. 14.7 a. Calculate the pH of 0.10 M HNO 3 b. Calculate the pH of 1.0 x 10 -10 M HCl

23. Weak Acids  Ka will be small.  ALWAYS WRITE THE MAJOR SPECIES.  It will be an equilibrium problem from the start.  Determine whether most of the H + will come from the acid or the water.  Compare Ka or Kw  The rest is just like previous equilibrium problems

24. Sample Ex. 14.8  The pH of Weak Acids  The hypochlorite ion (OCl - ) is a strong oxidizing agent often found in household bleaches and disinfectants. It is also the active ingredient that forms when swimming pool water is treated with clorine. In addition to its oxidizing abilities, the hypochorite ion has a relatively high affinity for protons (it is a much stronger base than Cl -, for example) and forms the weakly acidic hypochlorous acid (HOCl, Ka = 3.5 x 10 -8 ).  Calculate the pH of a 0.100 M aqueous solution of hypochlorous acid.

25. A mixture of Weak Acids  The process is the same.  Determine the major species.  The stronger will predominate.  Bigger Ka if concentrations are comparable  Sample Ex. 14.9: Calculate the pH of a mixture 1.00 M HCN (Ka = 6.2 x 10 -10 ) and 5.00 M HNO 2 (Ka = 4.0 x 10 -4 )  Calculate the concentration of cyanide ion (CN-)

26. Bases The OH - is a strong base. 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. Used as antacids because [OH - ] can’t build up.

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

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

29. Sample Ex. 14.13 The pH of Weak Bases Calculate the pH for a 15.0 M solution of NH 3 (K b = 1.8 x 10 -5 )

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

31. Polyprotic acid 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 Base in first step is acid in second. In calculations we can normally ignore the second dissociation.

32. Sulfuric acid is special In first step it is a strong acid. Ka 2 = 1.2 x 10 -2 Calculate the concentrations in a 1.0 M solution of H 2 SO 4 (Ex. 14.16) Calculate the concentrations in a 1.0 x 10 -2 M solution of H 2 SO 4 (Ex. 14.17)

33. Structure and Acid base Properties 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 more polar the X-H bond the stronger the acid (use electronegativities). The more polar H-O-X bond -stronger acid.

34. Strength of Carboxylic Acid The strength of a carboxylic acid depends upon the stability of the anion that it forms when it loses its labile (mobile proton). All carboxylic acids form this equilibrium, where ‘R’ can vary. RCOOH   RCOO- + H+ If the anion (RCOO-) is relatively stable tends to donate H+ ions and the acid will be relatively strong

35. The table below indicates that the methanoate ion is more stable than the ethanoate ion, which is more stable than the propanoate ion. (the smaller the pKa the stronger the acid)

36. The ethanoate and propanoate ions have electron- releasing alkyl groups (CH 3 and C 2 H 5 ) that ‘pump’ electrons into the COO- anion and make it reactive and unstable. As a result, the H+ ions tend not to be released. This ‘pumping’ of electrons is called the inductive effect. Methanoic acid on the other hand does not have an alkyl group so it does not push electrons into the COO–, the anion remains relatively stable.

37. A similar effect is observed when comparing the relative strengths of halogen substituted carboxylic acids. It is found that as the number of chlorine atoms increases, so does the strength of the acid.

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

39. Strength of oxyacids Electron Density ClOH

40. Strength of oxyacids Electron Density ClOHO

41. Strength of oxyacids ClOH O O Electron Density

42. Strength of oxyacids ClOH O O O Electron Density

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

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