Experiment7

Experiment7 - CHEM 241L: Preparation and Investigation of...

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CHEM 241L: Preparation and Investigation of Buffers 137 7 Experiment 7 Buffer Lab: Preparation and Investigation of Buffers Read Sections 8-1, 8-2, 8-3 (Chapter 8), 9-5 (Chapter 9), and Section 10-2 (Chapter 10) in Harris, 7 th ed. Note: You need to bring your laptop to lab for this experiment. PRE-LAB QUESTIONS The pre-lab questions are part of your lab report grade. You may not receive ANY outside help to answer the below two questions. You may use your Harris textbook and 241 notes but that is it. 1. You need to prepare 100 mL of a pH 5.00 acetate buffer that is 0.200 M in acetic acid (CH 3 CO 2 H; FW 60.05). You have pure liquid acetic acid and pure sodium acetate (CH 3 CO 2 - Na + ) available to you to prepare the buffer. The pKa of acetic acid is 4.756 and the density of liquid acetic acid is 1.049 g/mL. a. Calculate the volume (in mL) of liquid acetic acid that must be added to the 100 mL volumetric flask. b. Calculate the mass (in grams) of sodium acetate that you would need to add to the 100 mL flask to prepare 100 mL of a pH 5.00 acetate buffer. 2. You want to prepare 100 mL of a pH 5.00 acetate buffer that is 0.200 M in acetic acid, BUT you only have pure liquid acetic acid and 1.0 M NaOH available to you to prepare the buffer. What volume (in mL) of 1.0 M NaOH that must be added to the volume (in mL) of liquid acetic acid (Problem 1, Part a above) to give a buffer with a pH of 5.00. BACKGROUND & THEORY Buffered Solutions Buffered solutions are solutions that show very small changes in pH when small amounts of strong acid are added, when small amounts of strong base are added, or when the solution is diluted. Buffered solutions play an important in a vast array of chemical and biochemical systems. The pH of blood is primarily regulated by the carbon dioxide-carbonic acid-bicarbonate buffer Reading Assignment
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CHEM 241L: Preparation and Investigation of Buffers 138 system, which maintains a constant blood pH of 7.40. 1 Chemically, a buffered solution contains appreciable amounts of both a weak acid and its conjugate base. For instance, if you add 3 moles of formic acid (HCO 2 H) and 2 moles of formate (HCO 2 - ) to a 1.0 L flask and dilute to the mark, you have a buffer because there is enough weak acid and conjugate base to be capable of equally reacting with either added acid or base. If you simply added 3 moles of formic acid to 1.0 L of water, the acid will dissociate to a small extent producing a very small amount of conjugate base (formate). HCO 2 H + H 2 O ' HCO 2 - + H 3 O + K a = 1.80 x 10 -4 This is not a buffer because there is too little conjugate base present in solution. If you added a little strong acid to this solution there would not be enough conjugate base present to consume the added acid and thus resist a change in pH. If you simply added 3 moles of formate (a weak base) to 1.0 L of water, the base will
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This note was uploaded on 01/03/2009 for the course CHEM 241 taught by Professor Tiani during the Spring '08 term at UNC.

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Experiment7 - CHEM 241L: Preparation and Investigation of...

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