1 / 23

Buffers

Buffers. A buffer is a solution that is highly resistant to changes in pH when a strong acid or base is added. A buffer solution also has a pH close to the pKa of the weak acid in the buffer system. Buffer Systems.

barrett
Download Presentation

Buffers

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Buffers • A buffer is a solution that is highly resistant to changes in pH when a strong acid or base is added. • A buffer solution also has a pH close to the pKa of the weak acid in the buffer system

  2. Buffer Systems • Buffers are used to maintain a constant pH and are widely used in commercial products and a variety of lab procedures. • Buffers are made by mixing a weak acid with a solution containing its conjugate base. • When a strong acid is added the H+ ions are consumed by the conjugate base in the buffer. • When a strong base is added the OH- ions are consumed by the conjugate acid in the buffer.

  3. How It Works

  4. Without A Buffer A: A 100-mL sample of dilute HCl is adjusted to pH 5.00. B: After the addition of 1 mL of 1 M HCl (left) or of 1 M NaOH (right), the pH changes by several units.

  5. With A Buffer A: A 100-mL sample of a buffered solution, made by mixing 1 M CH3COOH with 1 M CH3COONa, is adjusted to pH 5.00. B: After the addition of 1 mL of 1 M HCl (left) or of 1 M NaOH (right), the pH change is negligibly small.

  6. Henderson-Hasselbalch • This equation is used to calculate the pH of a buffer: • Or the following equation can be used to calculate the [H+]: [H+] = Ka x nHB nB-

  7. Calculating pH • Lactic acid, C2H5OCOOH, is a weak organic acid present in both sour milk and buttermilk. A buffer is prepared by dissolving 1.00 mol of lactic acid (Ka = 1.4 x 10-4) and 1.00 mol of sodium lactate in enough water to form 550mL of solution. Calculate the pH of the buffer. • Ans: 3.85

  8. Calculating pH • Calculate the pH: (a) Of a buffer solution consisting of 0.50 M CH3COOH and 0.50 M CH3COONa. Ans: 4.74 (b) After adding 0.020 mol of solid NaOH to 1.0 L of the buffer solution in part (a). Ans: 4.77 (c) After adding 0.020 mol of HCl to 1.0 L of the buffer solution in part (a). Ans: 4.70

  9. Buffer Capacity • It follows that a buffer has the highest capacity when the component concentrations are equal, that is, when [A−]/[HA] = 1: • For a given concentration, a buffer whose pH is equal to or near the pKa of its acid component has the highest buffer capacity:

  10. Preparing a Buffer Solution • 1. Choose the conjugate acid-base pair: First, decide on the chemical composition of the buffer, that is, the conjugate acid-base pair. This choice is determined to a large extent by the desired pH. Remember that a buffer is most effective when the ratio of its component concentrations is close to 1, in which case the pH ≈ pKa of the acid. • Reference Table14.1 – page 382 • *know the acetic acid and ammonia system vlaues

  11. Preparing a Buffer (cont.) • 2. Calculate the ratio of buffer component concentrations: Next, find the ratio of [A−]/[HA] that gives the desired pH.

  12. Preparing a Buffer (cont.) • 3. Determine the buffer concentration: Next, decide how concentrated the buffer should be. Remember that the higher the concentrations of components, the greater the buffer capacity. For most laboratory-scale applications, concentrations of about 0.5 M are suitable, but the decision is often based on availability of stock solutions.

  13. Preparing a Buffer (cont.) • 4. Mix the solution and adjust the pH: Thoroughly dissolve the calculated amount of salt containing the moles of conjugate base desired in the desired concentration of the weak acid to a final volume of 1 L. Finally, note that because of the behavior of nonideal solutions, a buffer prepared in this way may vary from the desired pH by a few tenths of a pH unit. Therefore, after making up the solution, adjust the buffer pH to the desired value by adding strong acid or strong base, while monitoring the solution with a pH meter.

  14. Preparing a Buffer • Suppose you need to prepare a buffer with a pH of 9.00. • Which of the following buffer systems would you choose? (hint: choose the acid with a pKa=pH you want) • Ans: ammonium-ammonia system • What should be the ratio of the concentration of weak acid, HB, to its conjugate base, B- ? • Ans: 1.8 • What mass in grams of B- should be added to 100.0 mL of 1.00 M HB to give a pH of 9.00? • Ans: 0.95 g NH3

  15. Indicators in Acid-Base Titrations Most have a range of about 2 pH units, in keeping with the useful buffer range of 2 pH units (pKa ± 1). (pH range depends to some extent on the solvent used to prepare the indicator.)

  16. The acidic form of bromthymol blue is yellow (left) and the basic form is blue (right). Over the pH range in which the indicator is changing, both forms are present, so the mixture appears greenish (center).

  17. Strong Acid-Strong Base

  18. Weak Acid-Strong Base

  19. Weak Base-Strong Acid

  20. FRQ Practice #1 Hypochlorous acid, HOCl, is a weak acid commonly used as a bleaching agent. The acid–dissociation constant, Ka, for the reaction represented below is 3.2×10–8. HOCl ↔ OCl– + H+ (a) Calculate the [H+] of a 0.14–molar solution of HOCl.

  21. FRQ Practice #1 (continued) (b) Write the correctly balanced net ionic equation for the reaction that occurs when NaOCl is dissolved in water and calculate the numerical value of the equilibrium constant for the reaction. (c) Calculate the pH of a solution made by combining 40.0 milliliters of 0.14–molar HOCl and 10.0 milliliters of 0.56–molar NaOH. And explain why Ms. Andre is so cool. (d) How many millimoles of solid NaOH must be added to 50.0 milliliters of 0.20–molar HOCl to obtain a buffer solution that has a pH of 7.49? Assume that the addition of the solid NaOH results in a negligible change in volume.

  22. FRQ Practice #2 Methylamine, CH3NH2, is a weak base that reacts according to the equation below. The value of the ionization constant, Kb, is 5.2510–4. Methylamine forms salts such as methylammonium nitrate, (CH3NH3+)(NO3–). CH3NH2 + H2O ↔ CH3NH3+ + OH– (a) Calculate the hydroxide ion concentration, [OH–], of a 0.225–molar aqueous solution of methylamine.

  23. FRQ Practice #2 (continued) (b) Calculate the pH of a solution made by adding 0.0100 mole of solid methylammonium nitrate to 120.0 milliliters of a 0.225–molar solution of methylamine. Assume no volume change occurs. (c) How many moles of either NaOH or HCl (state clearly which you choose) should be added to the solution in (b) to produce a solution that has a pH of 11.00? Assume that no volume change occurs. (d) A volume of 100. milliliters of distilled water is added to the solution in (c). How is the pH of the solution affected? Explain.

More Related