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Chapter 16 Acids and Bases

Chapter 16 Acids and Bases. Drill Determine which strong acid and strong base the following salts were derived from: LiCl Ba 3 (PO 4 ) 2 CaSO 4 Sr(NO 3 ) 2. Objectives SWBAT Distinguish between Arrhenius, Bronsted Lowry and Lewis acids and bases. Arrhenius Definition. Definitions:

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Chapter 16 Acids and Bases

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  1. Chapter 16 Acids and Bases

  2. Drill • Determine which strong acid and strong base the following salts were derived from: • LiCl • Ba3(PO4)2 • CaSO4 • Sr(NO3)2

  3. Objectives SWBAT Distinguish between Arrhenius, Bronsted Lowry and Lewis acids and bases.

  4. Arrhenius Definition • Definitions: • Acids produce hydrogen ions in aqueous solution. • Bases produce hydroxide ions when dissolved in water. • Limited to aqueous solutions. • Only one kind of base (hydroxide). • NH3 ammonia could not be an Arrhenius base.

  5. Bronsted-Lowry Definitions • Definition:An acid is a 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- • Water is a base that makes a hydronium ion

  6. Remember Strong acids completely dissociate in water. HCl + H2O  H3O +1 + Cl-1 This reaction goes to completion and there is no HCl left in the solution. Use a single direction arrow.

  7. Remember Weak acids only partially dissociate. CH3COOH + NH3⇔ CH3COO-1 + NH4+1 This is an equilibrium reaction. There are significant amounts of reactants and products in the solution. Use a double headed arrow. ⇔

  8. Remember Hydroxides (and some oxides) are strong bases. All other common bases are weak. Weak bases establish an equilibrium system like acids.

  9. Acid Base Pairs • General equation • HA(aq) + H2O(l) H3O+(aq) + A-(aq) • Acid + Base Conjugate acid + Conjugate base • This is an equilibrium situation. • There is competition for H+ between H2O and A- • The stronger base controls direction of the rxn. • If H2O is a stronger base it takes the H+ • Equilibrium would then move to right.

  10. Use the following reaction and the conjugate acid/base chart to determine which direction the equilibrium will lie. CH3COOH + NH3⇔ CH3COO-1 + NH4+1 CH3COOH is a stronger acid than NH4+1 NH3 is a stronger base than CH3COO-1 The equilibrium will favor the side in which the weaker acid and base a present. Equilibrium will lie to the right.

  11. Acid Dissociation Constant Ka • HA(aq) + H2O(l) H3O+(aq) + A -1(aq) • Ka = [H3O+1][A-1] [HA] • H3O+1 is often written H+1 ignoring the water in equation (it is implied). • Since this is the equilibrium constant associated with weak acid dissociation, this particular Kc is most commonly called the acid dissociation constant Ka

  12. 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 lies far to right. • Conjugate base must be weak.

  13. Strong acids Kais large [H+] is equal to [HA] A-1 is a weaker base than water Weak acids Kais small [H+] <<< [HA] A-1 is a stronger base than water Back to Pairs

  14. Types of Acids • Monoprotic Acids have only one hydrogen. • 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.

  15. Amphoteric • Amphoteric means that the substance can behave as both an acid and a base. • Water molecules interact with each other and ionize. At the same time, the ions in solution reform molecules of water as shown in the following reaction. (This means that water auto-ionizes) 2H2O(l) H3O+1(aq) + OH-1 (aq) KW = [H3O+][OH-] = [H+][OH-]

  16. In pure water the concentrations of H3O+1 and OH-1 will always be equal. • [H+] = [OH-] = 1.0 x 10-7 • At 25ºC KW = 1.0 x10-14 • Therefore: • Neutral solution [H+] = [OH-]= 1.0 x10-7 • Acidic solution [H+] > [OH-] • Basic solution [H+] < [OH-]

  17. pH • In 1909, Danish biochemist S. P. L Sorensen introduced the pH system. • pH representing power of hydrogen

  18. pH • pH= -log[H+] • Used because [H+] is usually very small • As pH decreases, [H+] increases exponentially • Other equations: • pOH= -log[OH-] • pKa = -log K

  19. Sig Figs for pH Sig figs: the number of sig figs in the lead number is the number of decimal places for the pH value. (only the digits after the decimal place of a pH are significant) [H+] = 1.0 x 10-8 pH= 8.00 2 sig figs

  20. Relationships • Derivation: • KW = [H+][OH-] • -log KW = -log([H+][OH-]) • -log KW = -log[H+]+ -log[OH-] • pKW = pH + pOH • 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.

  21. [H+] 100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14 pH 0 1 3 5 7 9 11 13 14 14 13 11 9 7 5 3 1 0 pOH 10-14 10-13 10-11 10-9 Basic 10-7 10-5 10-3 10-1 100 [OH-] Acidic Neutral Basic

  22. Strong Acids • HBr, HI, HCl, HNO3, H2SO4, HClO4 • These acids completely dissociate • Therefore, [H+] = [HA] • 10-14 = [H+][OH-]

  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 • Rest is just like equilibrium chapter.

  24. Example • Calculate the pH of 2.0 M acetic acid HC2H3O2 with a Ka 1.8 x10-5 • Calculate pOH, [OH-], [H+]

  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 • Calculate the pH of a mixture 1.20 M HF (Ka = 7.2 x 10-4) and 3.4 M HOC6H5 (Ka = 1.6 x 10-10)

  26. 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

  27. The Other Way • What is the Ka of a weak acid that is 8.1 % dissociated as 0.100 M solution?

  28. 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.

  29. 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]

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

  31. Polyprotic Acids • Always dissociate stepwise. • The first H+ comes of much easier than the second. • Ka for the first step is much bigger than Ka for the second. • Denoted Ka1, Ka2, Ka3

  32. Polyprotic Acids • What does K stand for? • Is it easier to remove the first or second ionizable proton? • Is is easier to remove the first. • The K values become successively smaller as successive protons are removed. • You will need to do 2 or more ice boxes.

  33. Polyprotic Acid H2CO3 H+ + HCO3-1 Ka1= 4.3 x 10-7 HCO3-1 H+ + CO3-2 Ka2= 4.3 x 10-10 • Base in first step is acid in second. • In calculations we can normally ignore the second dissociation.

  34. Calculate the Concentration • Of all the ions in a solution of 1.00 M Arsenic acid H3AsO4 • Ka1 = 5.0 x 10-3 • Ka2 = 8.0 x 10-8 • Ka3 = 6.0 x 10-10

  35. Sulfuric Acid is Special • In first step it is a strong acid. • Ka2 = 1.2 x 10-2 • Calculate the concentrations in a 2.0 M solution of H2SO4 • Calculate the concentrations in a 2.0 x 10-3 M solution of H2SO4

  36. Salts • A salt is an ionic compound formed by the reaction between an acid and a base. • Salts are strong electrolytes that completely dissociate into ions in water. • Salts of the cation of strong bases and the anion of strong acids are neutral. • for example NaCl, KNO3

  37. Basic Salts • If the anion of a salt is the conjugate base of a weak acid - basic solution. • 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-] [HF]

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

  39. Basic Salts • Ka x Kb = [HF][OH-] x [H+][F-] [F- ] [HF] • Ka x Kb =[OH-] [H+] • Ka x Kb = KW

  40. 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+

  41. Acidic Salts • 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 • Ka x Kb = KW • Calculate the pH of a solution of 0.40 M NH4Cl (the Kb of NH3 1.8 x 10-5). • Other acidic salts are those of highly charged metal ions.

  42. Anion of weak acid, cation of weak base • Ka > Kb acidic • Ka < Kb basic • Ka = Kb Neutral

  43. 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.

  44. 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 oxygens are electronegative • Pull electrons away from hydrogen

  45. Cl O H Strength of Oxyacids Electron Density

  46. Cl O H Strength of Oxyacids Electron Density O

  47. Cl O H Strength of Oxyacids Electron Density O O

  48. Cl O H Strength of Oxyacids Electron Density O O O

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

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