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Acid and Base Equilibrium

Acid and Base Equilibrium. Chapter 16 Brown LeMay. Basic Concepts. Acids – sour or tart taste, electrolytes, described by Arrhenius as sub that inc. the H + conc. Bases – bitter to the taste, slippery, electrolytes, describes by Arrhenius as sub that inc. the OH - conc.

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Acid and Base Equilibrium

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  1. Acid and Base Equilibrium Chapter 16 Brown LeMay

  2. Basic Concepts • Acids – sour or tart taste, electrolytes, described by Arrhenius as sub that inc. the H+ conc. • Bases – bitter to the taste, slippery, electrolytes, describes by Arrhenius as sub that inc. the OH- conc.

  3. Dissociation of Water • Pure water exists almost entirely of water molecules. It is essentially a non-electrolyte. • Water ionizes to a small extent – auto-ionization • The equilibrium expression is H20(l) <-> H+(aq) + OH-(aq) Kw = [H+] [OH-]

  4. Since the H+ ion does not exist alone in water Kw is often expressed Kw = [H3O+] [OH-] because water conc is constant it does not appear in the expression The proton in water

  5. Values for Kw and [H+] and [OH-] • Kw = [H3O+] [OH-] = 1.0 x 10-14 1.0 x 10-14 = [X] [X] 1.0 x 10-14 = X2 X = 1.0 X 10-7 = [H3O+] = [OH-]

  6. holds true for situations not involving water Acids donate protons Bases accept protons The Bronsted – Lowry definition

  7. HCl(g) + NH3(g) NH4+(g) + Cl-(aq)\ donates H+ accepts H+ conj base conj acid BL-Acid BL-Base • notice that the reaction doesn’t happen in water and that the OH- concentration has not increased

  8. Conjugate Acids & Bases and Amphoteric Substances HNO2(aq) + H2O(l) NO2-(aq)+H3O+(aq) conjugate acid base base acid donates - accepts protons • note HNO2 is considered a Arrenius acid H+ inc in H20 also a BL because donates a Proton • all Arrenius acids and bases are also BL acids and bases – however all BL acids and bases may or may not be arrenius acids or bases water is not an Arrenius base in this example

  9. every acid has a conjugate base formed from the removal of a proton from that acid every base has a conjugate acid formed from the addition of a proton to that base

  10. Amphotheric Substances • NH3(aq)+H20(l) HH4+(aq) + OH-(aq) Base Acid Conj Acid Conj Base p-acc p-donar note water is acting as an acid in this reaction and a base in the previous one that makes it a amphortic substance

  11. Strengths of Acid, Bases • The stronger the acid is the weaker its conjugate base (weaker acid  stronger conj base) • The stronger the base the weaker its conjugate acid (weaker base  stronger conj acid) • Stronger acids and bases ionize to a greater extent than do weak acids and bases.

  12. Strong Acids – dissociate completely into ions • HNO3(aq)  H+(aq) + NO3-(aq) the production of H+ ions from the acid dominates – ignore the H+ donated from the water it is insignificant • The  equilibrium lies so far to the right because HNO3 doesn’t reform • The neg log of the H+ from the acid determines pH. • Strong bases also dissociate completely and the conc of OH- from the base is the only factor considered when calculating pH.

  13. Strong Acids HClO4 prechloric H2SO4 sulfuric HI hydroiodic HCl hydochloric HBr hydrobromic HNO3 nitric Strong Bases GI hydroxides ex. NaOH,KOH G2 Hydroxides Sr(OH)2 GI Oxides ex. Na2O, K2O GI,II Amides ex. KNH2,Ca(NH2)2 Acids and Bases

  14. The pH scale • pH – is defined as the neg log (base-10) of the H+ ion concentration pH = -log [H+] • What is the pH of a neutral solution [H+] = 1.0 x 10^-7 pH = -log [1.0 x 10^-7 pH = 7

  15. Strong acids and the pH scale • An acidic solution must have a [H+] conc greater than 1.0 X 10-7 ex 1.0 X 10-6 • -log [1.0 X 10-6] = pH 6 • What is the pH of a basic solution? A basic solution is one in which the [OH-] is greater than 1X10-7.

  16. Calc. pH of strong basic solutions • Calculate the pH of a sol that has a [OH-] con. Of 1.0 X10-5 • Kw = 1x10-14 = [H+] [OH-] Kw = 1x10-14 = [H+] [1.0X10-5] [H+] = 1x10-14 = 1.0 X 10-9 1.0X10-5 pH = -log 1.0 x 10-9 pH = 9

  17. The “p” Scale • The negative log of a quantity is labeled p (quantity) • Ex: we could reference the quantity of OH- directly: pOH = -log[OH-] • From the definition of Kw -log Kw =(-log [H+]) (-log [OH-])= -log 1x10-4 Kw = pH + pOH = 14

  18. Calc. pH using the p scale • Ex. OH- conc = 1.0X10-5 • -log 1.0X10-5 = 5 = pOH • pH + pOH = 14 • pH + 5 = 14 • pH = 9

  19. Weak Acids • partially ionize in aqueous solution • mixture of ions and un-ionized acid in sol. • WA are in equilibrium (H20 is left out because it a pure liquid) HA(aq) + H20(l) H30+(aq) + A-(aq) • Ka is the acid dissociation constant • Ka = [H30] [A-] = [H+] [A-] [HA] [HA]

  20. The larger the Ka value the stronger the acid is – more product is in solution Acid Dissociation Constant

  21. Weak Bases • Weak bases in water react to release a hydroxide (OH-) ion and their conjugate acid: • Weak Base(aq) + H2O(l) Conjugate Acid(aq) + OH-(aq)

  22. A common weak base is ammonia • NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) • Since H2O is a pure liquid it is not expressed in the equilibrium Kb expression • Kb = [NH4+][OH-] (base dissociation [NH3] constant) • Kb always refers to the equilibrium in which a base reacts with H2O to form the conjugate acid and OH-

  23. Lewis Acids and Bases • Review - An Arrhenius acid reacts in water to release a proton - base reacts in water to release a hydroxide ion • In the Bronstead-Lowry description of acids and bases: acid reacts to donate a proton - a base accepts a proton

  24. G.N.Lewis defination - • Lewis acid is defined as an electron-pair acceptor • Lewis base is defined as an electron-pair donor

  25. In the example with ammonia, the ammonia is acting as a Lewis base (donates a pair of electrons), and the proton is a Lewis acid (accepts a pair of electrons)

  26. Lewis is consistent with the description by Arrhenius, and with the definition by Bronstead-Lowry. However, the Lewis description, a base is not restricted in donating its electrons to a proton, it can donate them to any molecule that can accept them.

  27. Calculating the pH of a Weak Acid • What is the pH of an aq sol that is 0.0030M pyruvic acid HC2H3P3? Ka = 1.4x10-4 at 25oC • HC2H3P3  H+ + C2H3P3- I 0.0030 0 0 C -X X X E 0.0030-X +X +X

  28. Ka = [H+] [C2H3P3-] pluggin in the values [HC2H3P3] from the table 1.4x10-4 = X2 (0.0030-X) 1.4x10-4 (0.0030) = X2 4.2x10-7-1.4x10-4x = X2 X2 + 1.4x10-4x-4.2x10-7 = 0 a quadratic ignore the neg sol x = 5.82 x10-4 pH = -log 5.82 x10-4 pH = 3.24

  29. Learning Check • What is the pH at 25oC of a solution made by dissolving a 5.00 grain tablet of aspirin (acetylsalicylic acid) in 0.500 liters of water? The tablet contains 0.325g of the acid HC9H7O4. Ka = 3.3x10-4 mm = 180.2g/l • H+ = 9.4x10-4 pH = 3.03

  30. Buffers • A solution that resists changes in pH when a limited amount of an acid or base is added to it. • Buffers contain either a weak acid and it’s conj. base or a weak base and it’s conj. acid.

  31. Examples • Ex. Weak acid and conj. base equal molar amounts • Strong Acid added H+ + A- HA conj. base weak acid the conj base interacts with the H+ ions from the strong acid changing them to a weak acid

  32. Strong base added OH + HA  HOH + A- weak acid conj base the weak acid interacts with the OH- ion from the base to form water and the conj. base If the concentration of A- and HA are large and the amount of H+ or OH- is small the solution will be buffered or the change in pH will be minimized.

  33. Buffering capacity – the amount of acid or base a buffer can react with before a significant change in pH occurs • Ratio of acid to conj base – unless the ratio is close to 1 ( between 1:10 and 10:1) will be too low to be useful.

  34. Calculating the pH of a buffer • Note: a solution of 0.10 M acedic acid and its conj base 0.20 M acetate from sodium acetate is a buffer solution pH = 5.07 • Ex. Calc. the pH of a buffer by mixing 60.0 ml of 0.100 M NH3 with 40.0ml of 0.100 M NH4Cl.

  35. 1St cal the conc. of each species M = moles/liters mol of NH3 0.10M = X/0.060 l = 0.0060mol mol of NH4 0.10M = X/0.040 l = 0.0040mol [NH3]= 0.0060 mol/ 0.100 l = 0.060 M [NH4] = 0.0040 mol/ 0.100 l = 0.040 M NH3 + H2O  NH4+ + OH- I 0.060M O.040 0 C -X +X +X E 0.060-x 0.040+x X

  36. Kb = [NH4+] [OH-] ( 0.040+X)X 1.8X10-5 [NH3] (0.060-X) Ignore X 1.8X10-5 = 0.040X X = 2.7X10-5 0.060 -log (2.7x10-5) = 4.6 pOH pH = 9.4 Or using Henderson - Hasselbalch equation pOH = pKb + log [conj acid] = 4.74 + log ( [0.04] [B] [0.06]) = 4.6 pOH pH= 9.4

  37. What is the pH of a buffer prepared by adding 30.0ml of .15M HC2H3O2 to 70ml of .2M NaC2H3O2? [HC2H3O2] = .15M = x/.03 = .0045/.01 = .045 [C2H3O2] = .20M =x/.07 = .0140/.01 = .140 ka = 1.7x10-5 HHeq pH = pKa + log [conj base] [acid] pH = 4.77 + log(.140 = 5.3 .045)

  38. Adding an acid or a base to a buffer • Calc the ph of 75ml of the buffer solution of(0.1M HC2H3O2 and 0.2M NaC2H3O2) to which 9.5 ml of 0.10M HCl has been added. Compare the change to that of adding HCl to pure water. H+ + C2H3O2 -  HC2H3O2 H+ = 0.10M = n/.0095l = 0.00095 moles C2H3O2 - = 0.2M = n/.075 = 0.0150 moles HC2H3O2 = 0.10M = n/.075 = 0.0075 moles

  39. Neutralization Reaction • C2H3O2 = C2H3O2 moles – H+ moles 0.0150n - .00095n = 0.014 • HC2H3O2= Orginial Conc. + Conc Contributed by reaction 0.075 moles + 0.00095 = 0.0085mol [C2H3O2]= 0.014mol/0.085l = 0.16M [HC2H3O2] = 0.0085/0.085 = 0.10M pH = 4.76 + log (.16/.10) = 4.96

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