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ACIDS AND BASES

ACIDS AND BASES. CHAPTER 15. I. Arrhenius Acids and Bases. (What we have been using to this point) Arrhenius Acid is a substance that, when dissolved in water, increases the concentration of hydrogen ion, H + or ( hydronium ion H 3 O + ).

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ACIDS AND BASES

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  1. ACIDS AND BASES CHAPTER 15

  2. I. Arrhenius Acids and Bases (What we have been using to this point) Arrhenius Acidis a substance that, when dissolved in water, increases the concentration ofhydrogen ion, H+ or (hydronium ion H3O+). Arrhenius Baseis a substance that, when dissolved in water, increases the concentration ofhydroxide ions, OH-.

  3. II. Bronsted-Lowry Acids & Bases A. Definitions 1. Bronsted Acid: H+ (proton) donor Ionizable hydrogen in the acid structure is usually bonded to an electronegative atom. 2. Bronsted Base: H+ (proton) acceptor Base structure must have an unshared pair of electrons.

  4. B. Broader Definition For Acid-Base Reactions Consider:HCl + NH4OH  NH4Cl + H2O Consider: HF + HCO3- H2CO3 + F-

  5. C. Amphiprotic Compounds Some compounds can act as a Bronsted acid in one case and a Bronsted base in another case. It depends on the molecules the amphiprotic compound is reacting with. Example: Water HCl + H2O  H3O+ + Cl- NH3 + H2O  NH4+ + OH-

  6. D. Conjugate Acid-Base Pairs A conjugate acid-base pair consists of two species in an acid-base reaction, one acid and one base, that differ by the loss or gain of a proton. Conjugate acid- species formed when a base has accepted a proton. Conjugate base- species formed when an acid has donated or removed a proton. Example: Acetic Acid CH3COOH ⇌CH3COO- + H+ CH3COOH + H2O⇌CH3COO- + H3O+

  7. Example Problems 1. What is the conjugate base of HNO3 ? 2. What is the conjugate acid of NH2- ? 3. Label all species in the following reaction. H2PO4- + HCO3-⇌H2CO3 + HPO42-

  8. E. Strong vs. Weak Acids and Bases 1. Strong Acids and Bases Essentially go to 100% ionization Rxn normally shown with single headed arrow. Strong AcidsStrong Bases HCl LiOH HBr NaOH HI KOH HNO3 Ca(OH)2 H2SO4 Ba(OH)2 HClO4 Sr(OH)2 HNO3 → H+ + NO3-

  9. 2. Weak Acids and Bases Partially dissociate or ionize, reaching some dynamic equilibrium state. Rxn normally shown with double headed arrow. HCN ⇌ H+ + CN- Weak acids and bases significant in biological systems.

  10. F. Examples of Organic Weak Acids and Bases Carboxylic Acids: (weak acids) RCO2H or RCOOH Anion formed by ionization is resonance stabilized Examples: pyruvic acid, lactic acid, acetic acid Amines: (weak bases) Based on ammonia, with differing # of “R” groups Primary, secondary, tertiary

  11. G. Polyprotic Acids and Bases Polyprotic acids can donate more than one proton sequentially. H2SO4 H3PO4 H2S Polyprotic bases can accept more than one proton sequentially. SO42- PO43- S2-

  12. III. Self-Ionization of Water Description Pure water slightly ionizes on its own (autoionizes) with equilibrium reactant favored. H2O(l) + H2O(l) ⇌H3O+(aq) + OH-(aq) Determine the Kc value.

  13. Kc for water ionization = Kw Kw = ionization constant for water or = ion product constant for water Kw = [H+] [OH-] = 1 x 10-14 at 25oC What are [H+] and [OH-] for pure water at 25oC?

  14. True for Water and Dilute Aqueous Solutions (Including Acid and Base Solutions of Interest) [H+] [OH-] = 1 x 10-14 or [H3O+] [OH-] = 1 x 10-14 Neutral Solution [H+] = [OH-] Acidic Solution [H+] > [OH-] Basic Solution[H+] < [OH-] or [H3O+]

  15. IV. pH Scale In most aqueous weak acid and base solutions, the H+ concentrations are very small. Easier to express concentrations using a logarithmic relationship. p = - log Thus:pH = -log [H+] pOH = -log [OH-]

  16. What is the relationship between pH and pOH? [H+] [OH-] = 1 x 10-14 (-log [H+]) + (-log [OH-]) = -log (1 x 10-14) pH + pOH = 14 pH + pOH = 14

  17. The pH Scale Since pH is a logarithmic scale, cola drinks (pH about 2.5) are about ____ times as acidic as tomatoes (pH about 4.5)

  18. Recall for problem solving: ****[H+] is the same as [H3O+] 1) pH = -log [H+] 2) [H+] = 10-pH orantilog (-pH) = [H+] 3) pOH = -log [OH-] 4) [OH-] = 10-pOH or antilog (-pOH-) = [OH-] 5) [H+] [OH-] = 1 x 10-14 6) pH + pOH = 14

  19. Problem Solving: 1) A blood sample (considered a dilute aqueous solution) contains 7.2 x 10-8 mol H+ per liter. a. What is the pH of the blood sample? b. What is the [OH-] of the blood sample? c. Is the blood sample acidic, basic, or neutral? 2) The pH of an aqueous solution is 8.7 a. What is the [H+]? b. What is the pOH?

  20. V. Strong Acids and Bases Strong acids and bases undergo essentially complete dissociation (ionization). Equilibrium expressions not needed for problem solving. Example: What is the pH and [H+] for a 0.05 M HCl solution?

  21. Problem Solving for Weak Acids and Bases 1. They only dissociate or ionize partially in solution. 2. This means we have an equilibrium situation, and can solve problems using the techniques we learned earlier for dealing with equilibrium problems. (ICE tables…) 3. Equilibrium Constants (Kc) Ka for weak acids Kb for weak bases

  22. VI. Weak Acid Equilibrium Rxns A. Acid Ionization Constant 1. Definition Give the acid ionization constant expression for the acid: HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq) or HA(aq) ⇌ H+(aq) + A-(aq) Ka = ?

  23. 2. What does the Ka tell you? Larger Ka stronger acid (more of the acid ionizes in water to form H+) See Table 15.5, for examples Which is the strongest acid? Acetic acid Ka = 1.8 x 10-5 Boric Acid Ka = 7.3 x 10-10

  24. B. Ka From Equilibrium Concentrations A 0.10 M solution of phenol (weak acid) has a pH of 5.43. Calculate the Ka for phenol.

  25. C. pH Calculations from Ka of Weak Acid The weak monoprotic acid, parahydroxybenzoic acid, has a Ka of 2.6 x 10-5. Answer the following questions for a 0.200 M solution of parahydroxybenzoic acid. 1. What is the pH of the solution? 2. What percent of the parahydroxybenzoic acid has ionized in this solution?

  26. D. Comment on Polyprotic Acids Example: H3PO4 Ka’s = 7.5 x 10-3, 6.2 x 10-8,3.6 x 10-13 1. Show the three acid ionization equations with their accompanying Ka values. 2. Which is the strongest acid? 3. How would you go about calculating the pH of a given H3PO4 solution?

  27. VII. Weak Base Equilibrium Rxns • Base Ionization Constant (base will accept proton from H2O, forming OH-) 1. Definition Give the base ionization constant expression for the base: B(aq) + H2O(l) ⇌ BH+(aq) + OH-(aq) Kb = ?

  28. 2. Relative Strength of Weak Bases a. The Kb for ammonia is 1.8 x 10-5. The Kb for phosphate ion is 2.8 x 10-2. Which is the strongest base? b. Which will have a higher pH: a 1.0 M ammonia solution, or a 1.0 M phosphate ion solution

  29. B. Kb From Equilibrium Concentrations The pain killer, morphine, is a weak base. A 0.01 M morphine solution has a pH of 10.1. Calculate the Kb for morphine.

  30. C. pH Calculations from Kb of Weak Base The weak base methylamine (CH3NH2) has a Kb of 5.0 x 10-4. Answer the following questions for a 0.080 M aqueous solution of methylamine. a. Write the chemical equation. b. Write the base ionization equilibrium expression. c. Calculate the [OH-], pOH, and pH for solution.

  31. VIII. Relationship of Ka and Kb A. Conjugate Acid and Base Reaction May Be Written in Both Directions. For Example: HA + H2O ⇌ H3O+ + A- Ka = A- + H2O ⇌HA + OH- Kb = ** Reaction will go in direction of stronger acid and base to the weaker acid and base. ** Compare Ka and Kb to decide direction.

  32. B. Mathematical Relationship of Ka and Kb (for previous equations) Thus: Ka x Kb = [H+] [OH-] = Kw Ka x Kb = 1 x 10-14 and pKa + pKb = 14

  33. C. Example Problems Consider the dihydrogen phosphate ion, H2PO41-. Ka = 6.2 x 10-8. 1. What is its conjugate base? 2. Show the chemical reaction to which Ka applies for H2PO41- 4. Determine the Kb value for HPO42-. 5. Which direction is the reaction favored?

  34. IX. Acid-Base Reactions of Salts(Ions as Acids and Bases) A. Salts 1. Ionic compounds (salts) formed from acid-base reaction. HA + MOH  M+A- + H2O acid base salt 2. The salt (MA) could leave the resulting aqueous solution to be neutral, acidic, or basic.

  35. B. Hydrolysis of Salts Salts may react with water (hydrolysis reactioninvolving the splitting of a water molecule) to produce acidic or basic solutions. For hypothetical ionic compound MA, consisting of Mx+ cation and Ax- anion, consider possible hydrolysis reactions and possible consequences: Mx+ + H2O  M(OH)(x-1)+ + H+acidic solution Ax- + H2O  HA(x+1)- + OH- basic solution If both cation and anion hydrolyze, compare Ka and Kb

  36. C. Predicting If Hydrolysis Occurs 1. Write the hypothetical hydrolysis rxns for salt. (one for cation and one for anion) 2. Look at the hypothetical hydrolysis products. 3. If a strong acid or strong base was hypothetically produced, these hydrolysis reactions DO NOT occur. No acid or base properties upon hydrolysis. Recall Table of Strong Acids and Bases 4. If a weak acid or base was produced, these hydrolysis reactions DO occur and produce acid or base properties. For hydrolysis products: H+ : acidic OH- : Basic

  37. D. Example Problems Will an aqueous solution of the salt be acidic, basic, or neutral? 1) NaCl 2) NH4Cl 3) NH4CN

  38. Cumulative Example Problem The cyanate ion, OCN-, is a weak base. The Ka for HOCN (cyanic acid) is 3.4 x 10-8. Calculate the pH of a 0.200 M solution of NaOCN.

  39. X. Lewis Acids and Bases A. Some Acid-Base Reactions Do Not Fit Either Arrehenius or Bronsted –Lowry Definitions. Consider: HCl + NaOH  NaCl + H2O Al3+ + 6 H2O  Al (H2O)63+

  40. B. Lewis Acids and Bases 1. Lewis Acid: electron pair acceptor (Accepts a pair of electrons to form a new bond) ** Usually has an incomplete octet of electrons 2. Lewis Base: electron pair donor (Donates a pair of electrons to form a new bond) ** Musthave at least one lone pair of electrons C. Identify as Lewis acid or Lewis base? H2O Cr3+ BF3

  41. XI. Acid Rain Read Section 15.12 What is acid rain and what are its environmental effects?

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