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Arrhenius Definition n Acids produce hydrogen ions in aqueous solution. n Bases produce hydroxide ions when dissolved in water. n Limits to aqueous solutions.

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Presentation on theme: "Arrhenius Definition n Acids produce hydrogen ions in aqueous solution. n Bases produce hydroxide ions when dissolved in water. n Limits to aqueous solutions."— Presentation transcript:

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2 Arrhenius Definition n Acids produce hydrogen ions in aqueous solution. n Bases produce hydroxide ions when dissolved in water. n Limits to aqueous solutions. n Only one kind of base. n NH 3 ammonia could not be an Arrhenius base.

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

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

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

6 Acid dissociation constant K a n HA(aq) H + (aq) + A - (aq) n K a = [H + ][A - ] [HA] n We can write the expression for any acid. n Strong acids dissociate completely. n Equilibrium far to right. n Conjugate base must be weak.

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

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

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

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

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

12 Problems n If a solution has a [H + ] of.0035M what is the pH? n [OH - ]? n pOH?

13 Problems n If a solution has a pOH of 9.28 what is the [H + ]? n If a solution has a pH of 9.28 what is the [OH - ]?

14 BasicAcidicNeutral [H + ] pH Basic [OH - ] pOH

15 Calculating pH of Solutions n Always write down the major ions in solution. n Remember these are equilibria. n Remember the chemistry. n Dont try to memorize there is no one way to do this.

16 Strong Acids n HCl, HBr, HI, HNO 3, H 2 SO 4, HClO 4 n Completely dissociated n [H + ] = [HA] n [OH - ] is going to be small because of equilibrium n = [H + ][OH - ] n If [HA]< water contributes H +

17 Weak Acids n Ka will be small. n It will be an equilibrium problem from the start. n Determine whether most of the H + will come from the acid or the water. n Compare K a or K w n Rest is just like last chapter.

18 Example n Calculate the pH of 1.6 M HCl(aq) n Calculate the pH of 1.6 x M HCl(aq) n Calculate the pH of 2.0 M acetic acid HC 2 H 3 O 2 with a Ka = 1.8 x10 -5 n Calculate pOH, [OH - ], [H + ]

19 A mixture of Weak Acids n The process is the same. n Determine the major species. n The stronger acid will predominate. n Bigger Ka if concentrations are comparable n Calculate the pH of a mixture 1.20 M HF (Ka = 7.2 x ) and 3.4 M HOC 6 H 5 (Ka = 1.6 x )

20 Percent dissociation n = amount dissociated x 100 initial concentration n For a weak acid percent dissociation increases as acid becomes more dilute. n Calculate the % dissociation of 1.00 M and M Acetic acid (Ka = 1.8 x ) n As [HA] 0 decreases [H + ] decreases but % dissociation increases. n Le Châtelier

21 The other way n What is the Ka of a weak acid that is 8.1 % dissociated as M solution?

22 Bases n The OH - is a strong base. n Hydroxides of the alkali metals are strong bases because they dissociate completely when dissolved. n The hydroxides of alkaline earths Ca(OH) 2 etc. are strong dibasic bases, but they dont dissolve well in water. n Used as antacids because [OH - ] cant build up.

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

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

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

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

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

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

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

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

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

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

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

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

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

36 K a tells us K b n The anion of a weak acid is a weak base. n Calculate the pH of a solution of 1.00 M NaF. Ka of HF is 7.2 x n H + + F - HF n The F - ion competes with OH - for the H +

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

38 Anion of weak acid, cation of weak base n K a > K b acidic n K a < K b basic n K a = K b Neutral n NH 4 CN –Ka for HCN is 6.2 x –K b for NH 3 is 1.8 x 10 -5

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

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

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

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

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

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

49 Lewis Acids and Bases n Boron triflouride wants more electrons. n 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

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