14.1 Arrhenius Definition Acids produce hydrogen ions (H+) in aqueous solution. Bases produce hydroxide ions (OH-) when dissolved in water. Limits to aqueous solutions. Only one kind of base. NH3 ammonia could not be an Arrhenius base.

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) + H2O(l) H3O+ + Cl- In this equilibrium, water is a base, makes hydronium ion.

(YDVD)

Conjugate Acid/Base Pairs HA(aq) + H2O(l) H3O+(aq) + A-(aq) Acid + Base Conjugate acid + Conjugate base Conjugate base: everything that remains of the acid molecule after a proton is lost. Conjugate acid: formed when the proton is transferred to the base. This is an equilibrium. Competition for H+ between H2O and A- The stronger base controls direction. If H2O is a stronger base it takes the H+ Equilibrium moves to right.

Practice Identify the acid, base, conjugate acid, conjugate base in the following reactions. #32 HONH3 + H2O HONH2- + H3O+ HOCl + C6H5NH2 OCl- + C6H5NH3+

Acid Dissociation Constant Ka The equilibrium constant for the general equation. HA(aq) + H2O(l) H3O+(aq) + A-(aq) Ka = [H3O+][A-] [HA] (YDVD) H3O+ is often written H+ ignoring the water in equation (it is implied).

Acid dissociation constant Ka HA(aq) H+(aq) + A-(aq) Ka = [H+][A-] [HA] We can write the expression for any acid. Strong acids dissociate completely. Equilibrium far to right. Conjugate base is therefore weak.

Practice Write the dissociation reaction and the corresponding Ka equilibrium expression for the following acids in water. #30 HC2H3O2 CH3NH3+ Co(H2O)63+

14.2 Acid Strength Strong acids Ka is large [H+] is equal to [HA] A- is a weaker base than water Weak acids Ka is small [H+] <<< [HA] A- is a stronger base than water Strong Acids: HBr, HCl, HI HNO3, HClO3, HClO4, H2SO4, HIO3

Weak Acids Its equilibrium lies far to the left. (CH3COOH) Yields a much stronger (it is relatively strong) conjugate base than water. (CH3COO) (YDVD)

Types of Acids Monoprotic - One acidic proton. 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.

Water as an Acid and a Base Wateris amphoteric - it behave as both an acid and a base. Water also autoionizes. (BDVD) 2H2O(l) H3O+(aq) + OH-(aq) KW= [H3O+][OH-] = [H+][OH-] At 25ºC KW = 1.0 x10-14 In EVERY aqueous solution. Neutral solution [H+] = [OH-] = 1.0 x10-7 Acidic solution [H+] > [OH-] Basic solution [H+] < [OH-]

Practice Order the following from strongest to weakest acid. H2O HCl HOC6H5 Ka = 1.6 x 10-10 HF Ka = 7.2 x 10-4

Practice #38 Which is the stronger base? Cl- or H2O H2O or NO2- CN- or OC6H5- HNO2 Ka = 4.0 x 10-4 HCN Ka = 6.2 x 10-10 HOC6H5 Ka = 1.6 x 10-10

Practice Calculate [H+] or [OH-] for each of the following solutions at 25°C and state whether the solution is neutral, acidic, or basic. Ex.14.3 1.0 x 10-5 M OH- 1.0 x 10-7 M OH- 10.0 M H+ At 60°C, Kw is 1 x 10-13 Ex.14.4 Using LeChatelier’s principle, predict whether the reaction 2H2O(l) H3O+(aq) + OH-(aq) is exothermic or endothermic. Calculate [H+] and [OH-] in a neutral solution at 60°C.

14.3 The pH Scale 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

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

100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 [H+] 1 3 5 7 9 11 13 14 pH Acidic Neutral Basic 1 3 5 7 9 11 13 14 pOH Basic 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 [OH-]

Practice Ex. 14.5 Calculate the pH and pOH for each solution at 25°C. A) 1.0 x 10-3 M OH- B) 1.0 M H+ C) 2.5 x 10-2 M H+

Practice Calculate the [H+] and [OH-] for each solution at 25°C. Identify each solution as acidic, basic or neutral. pH= 3.20 pH=11.2 pOH= 9.6

Practice The pH of ammonia was measured to be 11.89 at 25°C. Calculate the pOH, [H+], [OH-] for the sample. The pH of a sample of human blood was measured to be 7.41 at 25°C. Calculate the pOH, [H+], [OH-] for the sample. Ex. 14.6

14.4 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. Read p. 665. Very important!

Strong Acids 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+

Practice #49 A solution is prepared by mixing 50.0 mL of 0.050 M HCl and 1500.0 mL of 0.10 M HNO3. Water is added until the final volume is 1.00 L. Calculate [H+], [OH-] and the pH for this solution.

Practice #50 (WA) A solution is prepared by mixing 90.0 mL of 5.00 M HCl and 30.0 mL of 8.00 M HNO3. Water is added until the final volume is 1.00 L. Calculate [H+], [OH-] and the pH for this solution.

14.5 Weak Acids Ka will be small. List major species in solution. Choose species that can produce H+ and write reactions. Based on K values, decide on dominant equilibrium. Write equilibrium expression for dominant equilibrium. List initial concentrations in dominant equilibrium.

Weak Acids Define change at equilibrium (as “x”). Write equilibrium concentrations in terms of x. Substitute equilibrium concentrations into equilibrium expression. Solve for x the “easy way.” Verify assumptions using 5% rule. Calculate [H+] and pH.

Practice Calculate the pH of 0.100 M hypochlorous acid, HOCl, with a Ka = 3.5 x10-8 Ex.14.8 Calculate the pH of 2.0 M acetic acid, HC2H3O2, with a Ka = 1.8 x10-5. What is the pOH, [OH-], and [H+]?

How about a quadratic? #57 Calculate the pH of 0.020 M HF solution, with a Ka = 7.2 x10-4. What is the pOH, [OH-], and [H+]?

A mixture of Weak Acids The process is the same. Determine the major species. The stronger will predominate. Bigger Ka if concentrations are comparable Calculate the pH of a mixture 1.00 M HCN (Ka = 6.2 x 10-10) and 5.00 M HNO2 (Ka = 4.0 x 10-4). What is the concentration of the CN- ion in this solution at equilibrium. EX. 14.9

Percent dissociation = amount dissociated x 100 initial concentration For a weak acid percent dissociation increases as acid becomes more dilute. Calculate the % dissociation of 1.00 M and .00100 M Acetic acid (Ka = 1.8 x 10-5 As [HA]0 decreases [H+] decreases but % dissociation increases. Le Chatelier

Calculating Ka What is the Ka of lactic acid (HC3H5O3) that is 3.7 % dissociated as 0.100 M solution? Ex. 14.11 A 0.15 M solution of a weak acid, is 3.0% dissociated. Calculate Ka #65 The pH of a 1.00 x 10-2 M solution of cyanic acid (HOCN) is 2.77 at 25°C. Calculate Ka for HOCN from this result. #67

Calculating Ka WA In a 0.100 M solution of HF, the percent dissociation is 8.1%. Calculate the Ka #66 A 0.050M solution of trichloroacetic acid (CCl3CO2H) is the same pH as 0.040M HClO4 solution. Caculate the Ka for trichloroacetic acid. #68

14.6 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.

Bases without OH- Bases are proton acceptors. NH3 + H2O NH4+ + OH- It is the lone pair on nitrogen that accepts the proton. Many weak bases contain N B(aq) + H2O(l) BH+(aq) + OH- (aq) Kb = [BH+][OH- ] [B]

Strength of Bases Hydroxides are strong. Calculate the pH of a 5.0 x 10-2 M NaOH solution Ex. 14.12 Calculate the pH, [OH] and pOH of a 4.0 x 10-4 M Ca(OH)2 solution. #78a Calculate the concentration of of an aqueous Sr(OH)2 solution that has a pH of 10.5 #82 Others are weak. Smaller Kb - weaker base. Which is the stronger base: NH3 or Cl- NH3 or C5H5N

Practice: weak base Calculate the pH of 15.0 M ammonia NH3 with a Kb = 1.8 x10-5 Ex. 14.13 Calculate the pH of 1.0 M methylamine with a Kb = 4.38 x10-4 Ex. 14.14 Calculate the [OH-], [H+], and pH of 0.20 M pyridine with a Kb = 1.7x10-9 #86b

14.7 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 Ka1, Ka2, Ka3

Polyprotic acid H2CO3 H+ + HCO3- Ka1= 4.3 x 10-7 HCO3- H+ + CO3-2 Ka2= 5.6 x 10-11 Base in first step is acid in second. In calculations we can normally ignore the second dissociation.

Calculate the Concentration Calculate the pH of a 5.00 M Phosphoric acid, H3PO4, solution and the equilibrium concentrations of each species. Ex. 14.15 Ka1 = 7.5 x 10-3 Ka2 = 6.2 x 10-8 Ka3 = 4.8 x 10-10

Sulfuric acid is special In the first step, it is a strong acid. In the second step, Ka2 = 1.2 x 10-2 For solutions more dilute than 1.0 M, the dissociation of HSO4- is important. Solving requires the quadratic equation. Calculate the pH in a 1.0 M solution of H2SO4 Ex. 14.16 Calculate the pH in a 1.0 x 10-2 M solution of H2SO4 Ex. 14.17

14.8 Salts as acids and bases Salts are ionic compounds. Salts of the cation of strong bases and the anion of strong acids are neutral. For example NaCl, KNO3 There is no equilibrium for strong acids and bases. We ignore the reverse reaction.

Basic Salts If the anion of a salt is the conjugate base of a weak acid, a basic solution will be generated. In an aqueous solution of NaF The major species are Na+, F-, and H2O F- + H2O HF + OH- Kb =[HF][OH-] [F- ] but Ka = [H+][F-] = 7.2 x 10-4 [HF]

Basic Salts Ka x Kb = [HF][OH-] x [H+][F-] [F- ] [HF]

Basic Salts Ka x Kb = [HF][OH-] x [H+][F-] [F- ] [HF]

Basic Salts Ka x Kb = [HF][OH-] x [H+][F-] [F- ] [HF]

Basic Salts Ka x Kb = [HF][OH-] x [H+][F-] [F- ] [HF] Ka x Kb =[OH-] [H+]

Basic Salts Ka x Kb = [HF][OH-] x [H+][F-] [F- ] [HF] Ka x Kb =[OH-] [H+] Ka x Kb = KW Calculate the pH of a 0.30 M NaF solution. The Ka for HF is 7.2 x 10-4 Ex. 14.18

Ka tells us Kb The anion of a weak acid is a weak base. Calculate the pH of a solution of 1.00 M NaCN. Ka of HCN is 6.2 x 10-10 The CN- ion competes with OH- for the H+ Which is the stronger acid NH4+ or CH3NH3+? (NH3 : Kb= 1.8x10-5 CH3NH2 :Kb=4.4x10-4) #102 Arrange the following solutions in order of most acidic to most basic. KOH, KCl, KCN, NH4Cl, HCl #99

Acidic salts A salt with the cation of a weak base and the anion of a strong acid will be acidic. The same development as bases leads to Ka x Kb = KW Calculate the pH of a solution of 0.10 M NH4Cl (the Kb of NH3 1.8 x 10-5). Ex. 14.19 Other acidic salts are those of highly charged metal ions.(Al3+) More on this later.

Salts with Acidic and Basic ions Ka > Kb acidic Ka < Kb basic Ka = Kb neutral Predict whether an aqueous solution of each of the following salts will be acidic, basic or neutral. EX.14.21 NH4C2H3O2 NH4CN Al2(SO4)3 #105 in Text.

14.9 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.

Strength of oxyacids The more oxygen hooked to the central atom, the more acidic the hydrogen. HClO4 > HClO3 > HClO2 > HClO Remember that the H is attached to an oxygen atom. The oxygen’s are electronegative Pull electrons away from hydrogen

Strength of oxyacids Electron Density Cl O H

Strength of oxyacids Electron Density O Cl O H

Strength of oxyacids Electron Density O Cl O H O

Strength of oxyacids Electron Density O O Cl O H O

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

14.10 Acid-Base Properties of Oxides Non-metal oxides dissolved in water can make acids. SO3 (g) + H2O(l) H2SO4(aq) Ionic oxides dissolve in water to produce bases. CaO(s) + H2O(l) Ca(OH)2(aq)

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

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

Lewis Acids and Bases F H F B N H F H Boron triflouride wants more electrons. BF3 is Lewis acid NH3 is a Lewis base. F H F B N H F H

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

Comparing the Acids Model Acid Base Arrhenius H+ producer OH- Producer B-Lowry H+ donor H+ Acceptor Lewis e- pair acceptor e- pair donor