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Chapter 16: Acid-Base Equilibria

Chapter 16: Acid-Base Equilibria. John Hnatow and Ketan Trivedi. PowerPoint by Amrita Raja. www.t2i2edu.com. Section 16.1: The Common Ion Effect.

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Chapter 16: Acid-Base Equilibria

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  1. Chapter 16: Acid-Base Equilibria John Hnatow and Ketan Trivedi PowerPoint by Amrita Raja www.t2i2edu.com

  2. Section 16.1: The Common Ion Effect • Here, we will discuss the acid-base properties of a solution with two solutes containing the same (i.e. a common) ion. This ion can be a cation or an anion. • Consider two solutes, NaF and HF (for HF, Ka = 7.2 x 10-4). NaF is a salt and it dissociates completely into ions when mixed with water. This is expressed by the reaction: • NaF(s) Na+(aq) + F-(aq) www.t2i2edu.com

  3. Section 16.1: The Common Ion Effect (cont.) • In the same solution, we have HF (aq), a weak acid. The ionization of the weak acid is expressed as: HF(aq) H+(aq) + F-(aq) Ka = 7.2 x 10-4 • Note: The fact that Ka > Kw indicates that in this solution, H+ is produced predominantly by the dissociation of HF. www.t2i2edu.com

  4. Section 16.1: The Common Ion Effect (cont.) • In this solution, F- is the common ion produced from NaF and HF. The presence of the common ion, F-, may have a significant effect on the equilibrium dissociation of the weak HF acid. Let us focus on the ionization reaction of the weak acid, HF. HF(aq) H+(aq) + F-(aq) www.t2i2edu.com

  5. Section 16.1: The Common Ion Effect (cont.) HF(aq) H+(aq) + F-(aq) What is the consequence of adding NaF to an equilibrium solution of HF? • Adding NaF to the HF solution increases [F-] and disrupts the equilibrium. • Recall Le Chatelier’s Principle: “When a disturbance is imposed on a reaction system at equilibrium, the equilibrium shifts in the direction that reduces the effects of the disturbance.” • Hence, increasing [F-] by addition of NaF results in a shift of the equilibrium to the left. In turn, this shift reduces [H+], and the pH. This is called the common ion effect. www.t2i2edu.com

  6. Section 16.1: The Common Ion Effect (cont.) • In summary, the common ion effect is the shift of an equilibrium caused by addition of a solute having an ion in common with the equilibrium system. • The common ion effect has two important roles: • 1) It allows one to change the pH of a solution. • 2) It allows one to change the solubility of salts that are slightly soluble in water. www.t2i2edu.com

  7. Section 16.2-16.3: Equilibrium Calculations for Acids with the Common Ion Effect • Consider a general solution of a weak acid, HA, and a salt, NaA. The salt dissociates completely into ions as expressed by the reaction: NaA(s) Na+(aq) + A-(aq) • Note: The complete dissociation of the salt into its ions means that if we are given 0.1 M NaA, then [Na+] = 0.1 M, and [A-] = 0.1 M. www.t2i2edu.com

  8. Section 16.2-16.3: Equilibrium Calculations for Acids with the Common Ion Effect (cont.) • The ionization reaction of the weak acid, HA, is expressed as: HA(aq) H+(aq) + A-(aq) • Recall that A- is the weak conjugate base of the weak acid HA. The equilibrium constant for the ionization reaction is expressed as: www.t2i2edu.com

  9. Section 16.2-16.3: Equilibrium Calculations for Acids with the Common Ion Effect (cont.) • Rearrange the equation: • Take the negative log of both sides: Note: - log [H+] = pH and – log Ka = pKa www.t2i2edu.com

  10. Section 16.2-16.3: Equilibrium Calculations for Acids with the Common Ion Effect (cont.) Note: Switching the numerator and the denominator reverses the sign. www.t2i2edu.com

  11. Section 16.2-16.3: Equilibrium Calculations for Acids with the Common Ion Effect (cont.) • A- is the weak conjugate base of the weak acid HA. • This equation is called the Henderson-Hasselbalch equation. www.t2i2edu.com

  12. Section 16.2-16.3: Equilibrium Calculations for Acids with the Common Ion Effect (cont.) • One can use this equation to calculate the pH for solutions of compounds having a common ion. • In problems that involve the common ion effect, the initial concentrations of the salt and the weak acidare usually given. • As long as the initial concentrations of the salt and weak acid are greater than 0.1 M, we can neglect the ionization of the acidand the hydrolysis of the salt (reaction of the salt with water). • Under such conditions, we can use the initial concentrations for the acid and the conjugate basein the Henderson-Hasselbalch equation. www.t2i2edu.com

  13. Section 16.4-16.5: Equilibrium Calculations for Bases with the Common Ion Effect • The common ion effect can also be applied to a solution containing a weak base and a salt of the base. • Consider a solution of a weak base, NH3, and a salt, NH4Cl. The salt, NH4Cl, dissociates completely in H2O according to the reaction: NH4Cl (s) NH4+ (aq) + Cl -(aq) • We know that NH4+ is an acid, and that it can donate a proton to water. Thus, NH4+ (aq) NH3 (aq) + H+ (aq) www.t2i2edu.com

  14. Section 16.4-16.5: Equilibrium Calculations for Bases with the Common Ion Effect (cont.) NH4+ (aq) NH3 (aq) + H+ (aq) • The equilibrium constant expression for this reaction is: • Rearrange this expression to solve for [H+]. www.t2i2edu.com

  15. Section 16.4-16.5: Equilibrium Calculations for Bases with the Common Ion Effect (cont.) NH4+ (aq) NH3 (aq) + H+ (aq) • Take the negative log of both sides. www.t2i2edu.com

  16. Section 16.4-16.5: Equilibrium Calculations for Bases with the Common Ion Effect (cont.) NH4+ (aq) NH3 (aq) + H+ (aq) • Recall that NH4+ is the weak conjugate acid of the weak base, NH3. www.t2i2edu.com

  17. Section 16.4-16.5: Equilibrium Calculations for Bases with the Common Ion Effect (cont.) • From here onwards, whether we consider a weak acid and its conjugate base or a weak baseand its conjugate acid, the Henderson-Hasselbalch equation will be written as: www.t2i2edu.com

  18. Section 16.6-16.7: Buffer Solutions Buffer solutions have the ability to resist changes in pH upon addition of small amounts of either acidor base. www.t2i2edu.com

  19. Section 16.6-16.7: Buffer Solutions (cont.) • Buffers are very important in chemical and physiological processes. • The pH in the human body varies from one fluid to another. • For example, the pH of blood is 7.4, whereas the pH of gastric juices in the stomach is about 1.5. • Blood can absorb the acids and bases produced in biological reactions without changing its pH. • Thus, blood is a buffer solution. www.t2i2edu.com

  20. Section 16.6-16.7: Buffer Solutions (cont.) • A buffer solution consists of a weak acid and its salt or a weak base and its salt. • In other words, a buffer solution follows the concepts of common ion solutions. www.t2i2edu.com

  21. Section 16.6-16.7: Buffer Solutions (cont.) • An example of an acidic buffer solution is a mixed solution of the weak acid, HF, and its salt, NaF. • Similarly, an example of a basic buffer solution is a mixed solution of the weakbase, NH3, and its salt, NH4Cl. • By choosing appropriate solutes and concentrations, a buffer solution can be created for any desired pH. www.t2i2edu.com

  22. Section 16.8-16.9: The Working of a Buffer Solution • Consider a buffer solution of a weak acid, HA, and its salt, NaA. The salt dissociates as: • NaA(s) Na+(aq) + A-(aq) • and the weak acid as: • HA(aq) H+(aq) + A-(aq) • When the buffer solution is at equilibrium, we can calculate the pH using the Henderson-Hasselbalch equation. www.t2i2edu.com

  23. Section 16.8-16.9: The Working of a Buffer Solution (cont.) • Now, let’s disturb this buffer solution by adding some NaOH. NaOH dissociates completely in solution: • NaOH(s) Na+(aq) + OH-(aq) • OH- is a strong base and accepts a proton from the acid. In solution, we have the acid, HA. Thus, OH- reacts with HA as: • OH-(aq) + HA (aq) H2O (l) + A-(aq) www.t2i2edu.com

  24. Section 16.8-16.9: The Working of a Buffer Solution (cont.) • Hence, adding a strong base to an acid buffer leads to the formation of more A-. Now, look at the equilibrium dissociation of the weak acid. • HA(aq) H+(aq) + A-(aq) • The equilibrium constant for this reaction is expressed as: www.t2i2edu.com

  25. Section 16.8-16.9: The Working of a Buffer Solution (cont.) HA(aq) H+(aq) + A-(aq) • According to Le Chatelier’s principle, increasing [A-] leads to a shift of the equilibrium to the left in the direction that consumes A-. Thus, the shift of the equilibrium results in an increase in [HA]. • This implies that [H+] changes very little. Hence, the pH of an acid buffer solution does not change much when a strong base is added. www.t2i2edu.com

  26. Section 16.10-16.11: Preparing a Buffer Solution with a Specific pH • In order to calculate the pH of a buffer solution, one uses the Henderson-Hasselbalch equation. • HA(aq) H+(aq) + A-(aq) • Here, A- is the base and HA is the acid. www.t2i2edu.com

  27. Section 16.10-16.11: Preparing a Buffer Solution with a Specific pH (cont.) • Recall, it is the ratio [A-]/[HA] that is the most resistant to change when either H+ or OH- is added to the buffer solution. The concentrations of A- and HA in the buffer solution are usually chosen in such a way that the ratio [A-]/[HA] is as close as possible to 1. • Thus, pH = pKa + log (1) • log (1) = 0 • Hence, pH = pKa • Thus, when preparing a buffer solution of a desired pH, one should choose a weak acid of pKavalues as close as possible to the desired pH. www.t2i2edu.com

  28. Section 16.12: Acid-Base Titrations and pH Curves • Titration is used to determine the amount of acid or base in a solution. • The solution being analyzed, called the analyte, is the acid or the base solution in an Erlenmyer flask or a beaker. • The titrant is the acid or the base solution in the buret. • The progress of an acid-base titration is often shown on a graph of the pH of the solution being analyzed vs. the amount of titrant added. • Such a graph is called a pH curve, or a titration curve. www.t2i2edu.com

  29. Section 16.12: Acid-Base Titrations and pH Curves (cont.) • In this section, we will consider three types of reactions: • a) Strongacid-strong base titrations • b) Weak acid-strong base titrations • c) Strong acid-weak base titrations www.t2i2edu.com

  30. Section 16.14-16.15: pH Curve for Strong Acid - Strong Base Titrations • The plot of “pH vs. Volume of Titrant Added”is called the Titration Curve, or the pH Curve. • In the previous example, the titrant was NaOH. • Hence, the pH curve for a strong acid-strong basetitration is: www.t2i2edu.com

  31. Section 16.14-16.15: pH Curve for Strong Acid - Strong Base Titrations (cont.) • In this curve, the equivalence point is at pH = 7.00. • Before the equivalence point, [H+] and pH can be calculated by dividing the moles of [H+] remaining by the total volume of the solution (acid+ base) in Liters. • After the equivalence point, [OH-] and the pOH can be calculated by dividing the moles of OH- remaining by the total volume of the solution (acid + base) in Liters. • Note: In order to calculate pH, use the relation pH + pOH = 14.00. www.t2i2edu.com

  32. Section 16.14-16.15: pH Curve for Strong Acid - Strong Base Titrations (cont.) • If the titrant is HCl, then the pH curve for a strong acid-strong base titration is: www.t2i2edu.com

  33. Section 16.14-16.15: pH Curve for Strong Acid - Strong Base Titrations (cont.) • In this curve, the equivalence point is at pH = 7.00. • Before the equivalence point, [OH-] and the pOH can be calculated by dividing the moles of OH- remaining by the total volume of the solution (acid + base) in Liters. • After the equivalence point, [H+] and the pH can be calculated by dividing the moles of H+ remaining by the total volume of the solution (base + acid) in Liters. • Note: In order to calculate pH, use the relation pH + pOH = 14.00. www.t2i2edu.com

  34. Section 16.16: Weak Acid - Strong Base Titrations • Calculating the pH during the titration of a weak acid by a strong base is a two-step process. • 1) The stoichiometric step: The reaction of OH- with a weak acid is assumed to be complete. The concentrations of acid remaining in solution and conjugate base formed are determined. • 2) The equilibrium step: The equilibrium constant for dissociation of the weak acid, the acid and the conjugate base concentrations are used to calculate [H+], then the pH. www.t2i2edu.com

  35. Section 16.17-16.18: pH Curve for Weak Acid - Strong Base Titrations • The plot of pH vs. Volume of Titrant added is called the Titration Curve, or pH Curve. The pH curve for a weak acid-strong base titration is shown to the right. • The equivalence point in an acid-base titration is defined by the stoichiometry. • Why? www.t2i2edu.com

  36. Section 16.17-16.18: pH Curve for Weak Acid - Strong Base Titrations (cont.) • Why? Because the equivalence point occurs when sufficient titrant has been added to react exactly with all the acid or base present. www.t2i2edu.com

  37. Section 16.17-16.18: pH Curve for Weak Acid - Strong Base Titrations (cont.) • In the titration curve of a weak acid by a strong base, the equivalence point is located at pH > 7.00. • At the equivalence point, the solution contains only a soluble salt and water. • The soluble salt consists of the cation from the strong base and the anion of the weak acid. • The former does not affect the pH, while the latter makes the solution basic (see Section in the Chapter on Acids and Bases). www.t2i2edu.com

  38. Section 16.17-16.18: pH Curve for Weak Acid - Strong Base Titrations (cont.) • The pH value at the equivalence point is affected by the strength of the acid. • The weaker the acid, the stronger its conjugate base and the higher the pH value at the equivalence point. www.t2i2edu.com

  39. Section 16.19: Strong Acid - Weak Base Titrations • Consider the titration of HCl, a strong acid, with NH3, a weak base. HCl ionizes as: • HCl(aq) H+(aq) + Cl-(aq) • NH3 is a base; thus, it reacts with the acid by accepting a proton. The reaction is: • NH3(aq) + H+(aq) NH4+(aq) www.t2i2edu.com

  40. Section 16.19: Strong Acid - Weak Base Titrations (cont.) • At the stoichiometric point, the solution contains H2O, NH4+, and Cl-. • Since NH4+ is a weak acid and Cl- is an extremely weak base, the solution is acidic at the stoichiometric point. • This means that NH4+ donates a proton to H2O, a weak base. Therefore, • NH4+(aq) + H2O (l) NH3(aq) + H3O+(aq) www.t2i2edu.com

  41. Section 16.19: Strong Acid - Weak Base Titrations (cont.) • Often, H2O is ignored and the reaction is written as: • NH4+(aq) NH3(aq) + H+(aq) • The important point to remember is that at the stoichiometric point, the solution is acidicand is characterized by Ka. • However, because we start off with a weak base (NH3) and the Kb value, we need to calculate Ka from Kb. www.t2i2edu.com

  42. Section 16.20-16.21: pH Curve for Strong Acid - Weak Base Titrations • The plot of “pH vs. Volume of Titrant Added”is called the Titration Curve or pH Curve. The pH curve for a strong acid-weak base titration is shown to the right. www.t2i2edu.com

  43. Section 16.20-16.21: pH Curve for Strong Acid - Weak Base Titrations (cont.) • In this curve, the equivalence point is located at pH < 7.00. This is because of the acidicnature of the conjugate acid of a weak base. • The equivalence point in an acid-base titration is defined by the stoichiometry. • In this case, the pH value at the equivalence point is determined by the strength of the base. • The weaker the base, the stronger its conjugate acid and the lower the pH value at the equivalence point. www.t2i2edu.com

  44. Section 16.22: Acid - Base Indicators • In acid-base titrations, the end point is observed by the change in color of an indicator. • An indicator is an organic dye whose color depends on the pH of the solution. • For example: The indicator methyl red is red at pH below 4. • This means that if the pH of a solution is 4 or below and a drop or two or methyl red indicator is added, then the solution turns red. www.t2i2edu.com

  45. Section 16.22: Acid - Base Indicators (cont.) • The indicator methyl red is yellow at pH above 7. • This means if the pH of a solution is 7 or higher and a drop or two of methyl red indicator is added, then the solution turns yellow. • As the pH changes from 4 to 7, the color of the indicator changes from red, to reddish-orange, to orange and to yellow. • Thus, through changes in color, an indicator displays the acidity, or the basicity of a solution. www.t2i2edu.com

  46. Section 16.22: Acid - Base Indicators (cont.) • Another very common acid-base indicator is phenolphthalein. • Phenolphthalein is colorless in solutions with a PH below 8. • Phenolphthalein turns bright pink when the pH of a solution is above 10. • Acid-base indicators are generally weak organic acids. • Hence, an indicator is often represented as HIn. HIn is a weak acidand its dissociation is represented as: • HIn (aq) H+(aq) + In-(aq) www.t2i2edu.com

  47. Section 16.22: Acid - Base Indicators (cont.) • Since HIn is a weak acid, it is characterized by its Ka value. Assume that the Ka value for the indicator HIn is 1.0 x 10-8. Thus, the equilibrium constant is expressed as: • Remember: HIn is a weak acid; thus, its conjugate base, In-, is a weak base. • Rearrange this equation as www.t2i2edu.com

  48. Section 16.22: Acid - Base Indicators (cont.) • Assume that we add one or two drops of indicator to a solution of pH = 1.0. If pH = 1.0, then [H+] = 0.10 M. www.t2i2edu.com

  49. Section 16.22: Acid - Base Indicators (cont.) • This means that the concentration of HIn is much greater than that of In-, and that the predominant species in solution is HIn. • Hence, if HIn was methyl red, we know that the indicator would turn the solution red since the pH is below 4. www.t2i2edu.com

  50. Section 16.22: Acid - Base Indicators (cont.) • Thus, a solution of pH = 1 with a few drops of methyl red appears red. • As a base is added, the concentration of H+ decreases. • This causes the equilibrium to shift to the right: • HIn (aq) H+(aq) + In-(aq) www.t2i2edu.com

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