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Chapter 15 Acid-Base Equilibria

CHEMISTRY. Chapter 15 Acid-Base Equilibria. Acids and Bases. Arrhenius’ Definition: Acids - are substances that produce hydrogen ions (protons or H + ) in solution. Bases - are substances that produce hydroxide ions in solution.

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Chapter 15 Acid-Base Equilibria

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  1. CHEMISTRY Chapter 15Acid-Base Equilibria

  2. Acids and Bases • Arrhenius’ Definition: • Acids - are substances that produce hydrogen ions (protons or H+) in solution. Bases - are substances that produce hydroxide ions in solution. Strong Acids and Strong Bases – totally ionize in solution Weak Acids and Weak Bases – partially ionize in solution

  3. Acid Dissociation in Water • General Rxn. when Acid dissolves in H2O • HCl + H2O H3O+ + Cl- acid base conj. Acid conj. base

  4. Properties of H2O @ 25 oC • H2O (l) DH+ (aq) + OH- (aq) • Neutral, can act as an acid and a base • Kw = [H+][OH-] = 1.0 x 10-14only @ 25 oC • Kw = water dissociation constant

  5. Acidity vs. Basicity • If [H+] >[OH-] , solution is acidic • If [H+] <[OH-] , solution is basic • The term pX = -log [concentration of X] • So: pH = -log [concentration of H+] • pOH = -log [concentration of OH-] • pH = power of hydrogen; the power of H to which 10 is raised

  6. pH • Kw = [H+][OH-] = 1.0 x 10-14 @ 25 oC • pKw = [-log H+] + [-log OH-]= - [log 1.0 x 10-14] = 14 • pH = -log [H+] • pOH = -log [OH-]

  7. Properties of H2O @ 25 oC • pKw = - [log 1.0 x 10-14] = 14 • pH + pOH = 14 • pH = 7 pOH = 7 • pH = pOH

  8. Things to Remember • pKw = - [log 1.0 x 10-14] = 14 @ 25 oC • pH + pOH = 14 • pH <7 ; acidic • pH > 7; basic • pH is between 0 - 14

  9. Broensted-Lowry’s Definition • Acid – is a proton (H+) donor. • Base – is a proton (H+) acceptor. • * Broensted-Lowry Definition is more general • It even applies to bases that have no –OH such as NH3.

  10. Terminologies • H+ = proton • OH- = hydroxide ion • H3O+ = hydronium ion • Conjugate base –acid minus proton • Conjugate acid – base plus proton

  11. More Terminologies • Conjugate acid-base pair • Consists of 2 substances related to each other by the donation and acceptance of a single proton (H+). • Acid Dissociation Constant (Ka)

  12. Equations • pH = - log [H+] • pOH = - log [OH-] • [H+] = 10 – pH • [OH-] = 10 - pOH • Kw = 10 - pKw • pKw = pH + pOH • pKw = - log [Kw] • Pw = [H+][OH-]

  13. Sample Problem • At 40 oC, a solution has Kw = 2.916 x 10-14; pH = 7.51 • Calculate the following: • A. pOH of the solution • B. hydrogen ion concentration [H+] • C. hydroxide ion concentration [OH-] • D. pKw • E. Is the solution acidic basic or neutral?

  14. Equilibrium • K = [H3O+ ][Cl- ] = [H+][Cl-] [HCl][H2O] [HCl] • H2O removed from top and bottom since H3O+ is simply H+ dissolved in water. • Remember: Keq = [products] [reactants]

  15. Acid Strength • Strength of acid is given by the equilibrium position of the dissociation reaction: • HA (aq) + H2O (l) H3O+ + A- • Strong acid – totally ionized and equilibrium lies far to the right • Weak acid – only partially ionized and equilibrium lies far to the left

  16. Strong Acid vs. Weak Acid • Strong Acid – yields a weak conjugate base (one that has weak affinity for proton; weaker than H2O) • Weak Acid – yields a strong conjugate base (one that has strong affinity for proton; stronger than H2O)

  17. Comparison

  18. Please Note! • Tuesday’s experiment is Experiment 29: Choice I.

  19. Sample Problems • Given [OH-] = 1.0 x 10-12 M, calculate pH. Is the solution basic, acidic or neutral? • Given [H+] = 4.30 x 10-6 M, calculate pH. Is the solution basic, acidic or neutral?

  20. Strong Acids and Bases • If the molarity of the acid or base is less than 10-6M then the autoionization of water needs to be taken into account. In other words, water is the primary source of H+ and OH-, so the pH would be neutral.

  21. Strong Acids and Bases • Strong Acids • The strongest common acids are HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4. • are strong electrolytes. • All strong acids ionize completely in solution:

  22. Strong Acids and Bases • If the molarity of the acid or base is less than 10-6M then the autoionization of water needs to be taken into account. In other words, water is the primary source of H+ and OH-, so the pH would be neutral.

  23. pH of Strong Acids and Bases • The pH (and hence pOH) of a strong acid is given by the initial molarity of the acid. • The pOH (and hence pH) of a strong base is given by the initial molarity of the base. • Be careful of stoichiometric ratios!

  24. Please Note! • Tuesday’s experiment is Experiment 29: Choice I.

  25. Bronsted-Lowry Acids and Bases • Bronsted-Lowry acids – compounds that donate a proton (H+) • Bronsted-Lowry Bases – compounds that accept a proton (H+) • Note that Bronsted-Lowry bases need not have the –OH group on the formula

  26. Weak Acids • Weak acids are only partially ionized in solution. • There is a mixture of ions and unionized acid in solution. • Therefore, weak acids are in equilibrium:

  27. NOTE • For Weak Acids and Weak Bases: • USE ICE to determine H+, OH-, pH and pOH.!

  28. Sample A Problem • A solution of 0.10 M formic acid (HCOOH) has a pH of 2.38 at 25 oC. • A. Calculate Ka for formic acid at this temperature. • B. What percent of this solution is ionized?

  29. Sample Problem • The Ka of acetic acid is 1.8 x 10-5. • A. Calculate the pH of a 0.30 M solution of CH3COOH. • B. Calculate OH- and pOH. • C. Calculate Kb. • Calculate % ionization.

  30. A Simple Trick • Use of approximation: eliminates the difficulty of quadratic equations. • Approximation is Valid if: X_______ x 100 < 5 % [Initial Concn.]

  31. Relationship between Ka and Kb • Ka x Kb = 1.0 x 10 -14only at 25 oC.

  32. pH of polyprotic acids • Treat polyprotic acids as separate steps! • #1. H2A (aq) D H+ (aq) + HA- (aq) Ka1 • # 2. HA- (aq) D H+ (aq) + A-2 (aq) Ka2 • Initial [H+] in Step 2 is Equil. [H+] from Step 1. • Total [H+] = SUM from Steps 1 & 2

  33. HOMEWORK • What is the pH of a 1.00 M solution of tartaric acid, H2C4H4O6 (aq.) at 25.0 oC? • Answer: pH = 1.49

  34. Sample Problem • The Ka of acetic acid is 1.8 x 10-5. Calculate the Kb of of CH3COOH.

  35. Sample Problem • The Ka of ammonia is 1.8 x 10-5. Calculate the pH of a 0.15 M solution of NH3.

  36. Sample Problem • Calculate the concentration of an aqueous solution of NaOH that has a pH of 11.50.

  37. HOMEWORK • What is the pH of a 1.00 M solution of tartaric acid, H2C4H4O6 (aq.) at 25.0 oC? • Answer: pH = 1.49

  38. Weak Acids • Calculating Ka from pH • Weak acids are simply equilibrium calculations. • The pH gives the equilibrium concentration of H+. • Using Ka, the concentration of H+ (and hence the pH) can be calculated. • Write the balanced chemical equation clearly showing the equilibrium. • Write the equilibrium expression. Find the value for Ka. • Write down the initial and equilibrium concentrations for everything except pure water. We usually assume that the change in concentration of H+ is x.

  39. Weak Acids • Calculating Ka from pH • Substitute into the equilibrium constant expression and solve. Remember to turn x into pH if necessary. • Using Ka to Calculate pH • Percent ionization is another method to assess acid strength. • For the reaction

  40. Sample Problem • A solution of NH3 in water has a pH of 10.50. What is the initial molarity of the solution?

  41. Other Weak Bases • Amines ex. Methylamine (CH3NH2) • carbonate ion (CO32-) • hypochlorite ion (ClO-1)

  42. Weak Bases • Also use ICE! • Calculation is the same as for weak acids! • Main difference is that you get [OH-] and pOH first.

  43. Effects of Salts on pH • Conjugate bases of strong acids have no effect on pH. • Conjugate acids of strong bases have no effect on pH. • Conjugate bases of weak acids increase pH (more basic). • Ex. F- (aq) + H2O(l) D HF (aq) + OH- (aq) • Conjugate acids of weak bases decrease pH (more acidic). • NH4+(aq) + H2O (l) D NH3 (aq) + H3O+ (aq)

  44. Relationship Between Ka and Kb

  45. Acid-Base Properties of Salt Solutions • Combined Effect of Cation and Anion in Solution • A cation that is the conjugate acid of a weak base will cause a decrease in the pH of the solution. • Metal ions will cause a decrease in pH except for the alkali metals (Grp. I) and alkaline earth metals.(Grp.II) • When a solution contains both cations and anions from weak acids and bases, use Ka and Kb to determine the final pH of the solution.

  46. Sample Problem • Determine whether the resulting solution in water will be acidic, basic or neutral. • A. K+ClO3- • B. Na+CH3COO- • C. Na2HPO4 Ka for HPO4- = 4.2 x 10-13 • D. NH4+Cl-

  47. Sample Problem • Predict whether the potassium salt of citric acid (K2+HC6H5O7-) will form an acidic, basic or neutral solution in water.

  48. Weak Acids • Polyprotic Acids • Polyprotic acids have more than one ionizable proton. • The protons are removed in steps not all at once: • It is always easier to remove the first proton in a polyprotic acid than the second. • Therefore, Ka1 > Ka2 > Ka3 etc.

  49. Weak Acids Polyprotic Acids

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