Chapter 15 Acid-Base Equilibria - PowerPoint PPT Presentation

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Chapter 15 Acid-Base Equilibria

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Chapter 15 Acid-Base Equilibria

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  1. Chapter 15Acid-Base Equilibria

  2. 15.1 – The common ion • Sometimes adding a salt adds an ion that was present in a weak acid. For example, look at what happens when you add NaF to a solution of HF… • The major species present are: Na+ F- H+ H2O

  3. 15.1 – The common ion • Because there is an excess of F- added to the acid, the equilibrium shifts to the reactants, creating more HF:

  4. 15.1 – The common ion • This is called the common ion effect and it is an application of Le Châtelier’s Principle If a change is made to the conditions of a chemical equilibrium, then the position of equilibrium will readjust so as to minimize the change made.

  5. 15.1 – The common ion • Calculations involving common ions: • Example,

  6. 15.2 – Buffered solutions • If you have enough of a weak acid and its conjugate base present (i.e., like we just saw), then you have a buffered solution – a solution which is capable of resisting changes in pH

  7. 15.2 – Buffered solutions • Buffers can also be prepared by using a weak base and its conjugate acid. For example, NH3 and NH4Cl form a buffered solution

  8. 15.2 – Buffered solutions • The Henderson-Hasselbach equation can be derived from the equilibrium expression for a weak acid, HX:

  9. 15.2 – Buffered solutions • The HH equation has many applications. • It relates the pH of a system to its pKa • Solving for [X-]/[HX] can give you an idea of the percent ionization of a particular compound • Can calculate the pH of a solution in an alternative method to the ICE diagram • Can aid in calculations where an acid/base has been added to buffered solutions

  10. 15.2 – Buffered solutions • Calculations involving buffered solutions • Example 1, • Calculate the pH of a 0.5L solution that is 1M HNO2 (Ka = 4.0 x 10-4) • Calculate the pH of a 0.5L solution that is 1M HNO2(Ka = 4.0 x 10-4) and 1M NaNO2

  11. 15.2 – Buffered solutions • Calculations involving buffered solutions • Example 1, • Calculate the pH of the solution in 1b if 0.010 mol of solid NaOH are dropped into it • Calculate the pH of an unbuffered solution of pH 7 when 0.010 mol of solid NaOH are dropped into it. Compare this to your answer in part c.

  12. 15.2 – Buffered solutions • Calculations involving buffered solutions • Example 1, • Calculate the pH of the solution in 1b if you added 1ml of concentrated (12M) HCl to it.

  13. 15.2 – Buffered solutions • Buffers are used extensively in biochemistry. They are found almost everywhere in biological systems. • The biological buffer worth noting is that of blood, composed of carbonic acid and carbonate ions

  14. 15.2 – Buffered solutions • A buffer’s range is the range of pH where it is most effective at buffering. It is usually within +/- 1 pH units of the pKa. • Example, the range of a acetic acid/acetate buffer is 3.74 – 5.74 because the pKa of acetic acid is 4.74.

  15. 15.3 – Buffering capacity • The buffering capacity of a buffered solution represents how much hydroxide/proton it can absorb without a significant change in pH • It is not given a value, but instead referred to as high/low, or shown in comparison with other solutions.

  16. 15.3 – Buffering capacity • An unbalanced buffer is a buffer that is composed of more of one component than the other. • For example, 0.1M CH3COOH/1M CH3COONa • This buffer is more effective at neutralizing an incoming acid than an incoming base (its capacity for neutralizing acid is higher than its capacity for neutralizing base)

  17. 15.3 – Buffering capacity • Selecting an ideal buffer • A buffer is most effective at a pH that is near the pKa of the weak acid. For example,

  18. 15.4 – Titrations and pH curves • A titration is a procedure performed to determine the concentration of acid or base in a solution. • An indicator (usually phenolphthalein) is often used to signify the endpoint of the titration. It indicates the equivalence point of a titration.

  19. 15.4 – Titrations and pH curves • Recall the terminology of titrations: • Titrant (known), in the burette • Anylate (unknown) in the E-flask

  20. 15.4 – Titrations and pH curves • A pH curve (or titration curve) is a graph with: • Volume of titrant added, on the x-axis • pH, on the y-axis

  21. 15.4 – Titrations and pH curves • There are 8 possibilities for a titration curve: AnylateTitrant SA SB SA WB WA SB WA WB SB SA SB WA WB SA WB WA Each one of these has a similar shape, although depending on the acid/base, they will be slightly different

  22. 15.4 – Titrations and pH curves • There are 3 important points or regions in a titration curve: • 1. Buffering region. In this region, the incoming base (or acid) is being neutralized by the excess of acid (or base) in the anylate. This creates small amounts of conjugate base (or acid), which makes a buffer. The result is that the pH changes very slowly.

  23. 15.4 – Titrations and pH curves • There are 3 important points or regions in a titration curve: • 2. Equivalence point. At this point, the number of moles of base added is equal to the number of hydrogen ions originally present (and vice versa)

  24. 15.4 – Titrations and pH curves • There are 3 important points or regions in a titration curve: • When you titrate a strong acid with a strong base (or vice versa), pH = 7 • When you titrate a weak acid with a strong base, pH > 7 • When you titrate a weak base with a strong acid, pH < 7

  25. 15.4 – Titrations and pH curves • There are 3 important points or regions in a titration curve: • 3. Half-equivalence point. At this point, [X-] = [HX], therefore pH = pKa. (Note, this does not apply to strong acid/base titrations. So you cannot determine the pKa of a strong acid using this method)

  26. 15.4 – Titrations and pH curves • Examples

  27. 15.4 – Titrations and pH curves • Examples

  28. 15.4 – Titrations and pH curves • Additionally, polyprotic acids, polybasic bases, and amphoteric substances have double curves.

  29. 15.4 – Titrations and pH curves • Amino acids have the general structure: • In humans, “R” is one of 20 different branches

  30. 15.4 – Titrations and pH curves • If protons are added to the environment, the amino group gets protonated, forming a zwitterion, a molecule with a negative charge on one side and a positive charge on the other (it is amphoteric): + H+ ⇌ “Zwitterion”

  31. 15.4 – Titrations and pH curves • If more protons are added, the zwitterion gets protonated, this time on the carboxylic acid side: + H+ ⇌ +H+ ⇌ “Zwitterion”

  32. 15.4 – Titrations and pH curves • There are Ka’s associated with each of these equilibria. If you dissolve an amino acid in an acid and titrate it with a base (i.e., NaOH), you can generate a titration curve where you can read the pKa of each of these + H+ ⇌ +H+ ⇌ “Zwitterion”

  33. 15.4 – Titrations and pH curves • In glycine, R = hydrogen • The isoelectric point, pI, is the pH where the net charge on the molecule is 0. The only molecule present at this point is the zwitterion. It is calculated as the average of pK’s:

  34. 15.5 – Acid-Base indicators • pH indicators are used extensively in titrationsto determine the endpoint. • An indicator is really just a weak acid. As you add base (or acid) to it, the ratio of acid to conjugate base changes. One is colorful, the other is not.

  35. 15.5 – Acid-Base indicators • Example, phenolphthalein: • A weak acid with Ka = 1.0 x 10-8 Summary of characteristics of phenolphthalein: Protonated, acidic form Deprotonated, basic form Abbreviated HIn Abbreviated In- Colorless Pink

  36. 15.5 – Acid-Base indicators • In order for our eyes to recognize the color, [In-] = 1 [HIn] 10 … So every indicator has a range of +/- 1 pKa in which it is useful.

  37. 15.5 – Acid-Base indicators • Choosing an indicator • Goal: get the endpoint and EP as close as possible to one another • Try to find an indicator whose range falls within the steep portion of the curve. Ideally, one whose pKa is equal to the equivalence point of the titration.

  38. 15.5 – Acid-Base indicators • Example, what indicator would you use in the titrationshown below? What color would the indicator change as you pass by the EP?

  39. 15.5 – Acid-Base indicators • You are titrating 50ml of 0.1M ammonia with 0.1M hydrochloric acid. (Kb = 1.8 x 10-5) • What will be the pH at the equivalence point? • What indicator should you use for this titration?