174 Experiment 6 Buffers 4. In this experiment, you will two buffer solutions. For each one, calculate the molar- prepare ity of the acid and the salt (conjugate base) in the solution. Then use the Henderson- Hasselbalch equation to calculate the expected initial molarity of each buffer. Refer to the worked example in the Introduction. a) Buffer #1: 20.0 mL of 0.200 M acetic acid solution and 10.0 mL of 0.200 M NAOH solution. n and 15.0 mL of 0 200 M NaOH

174 Experiment 6 Buffers 4. In this experiment, you will two buffer solutions. For each one, calculate the molar- prepare ity of the acid and the salt (conjugate base) in the solution. Then use the Henderson- Hasselbalch equation to calculate the expected initial molarity of each buffer. Refer to the worked example in the Introduction. a) Buffer #1: 20.0 mL of 0.200 M acetic acid solution and 10.0 mL of 0.200 M NAOH solution. n and 15.0 mL of 0 200 M NaOH

Chemistry: An Atoms First Approach

2nd Edition

ISBN:9781305079243

Author:Steven S. Zumdahl, Susan A. Zumdahl

Publisher:Steven S. Zumdahl, Susan A. Zumdahl

Chapter14: Acid- Base Equilibria

Section: Chapter Questions

Problem 4RQ: A good buffer generally contains relatively equal concentrations of weak acid and conjugate base. If...

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Fifty cm3 of 1.000 M nitrous acid is titrated with 0.850 M NaOH. What is the pH of the solution (a) before any NaOH is added? (b) at half-neutralization? (c) at the equivalence point? (d) when 0.10 mL less than the volume of NaOH to reach the equivalence point is added? (e) when 0.10 mL more than the volume of NaOH to reach the equivalence point is added? (f) Use your data to construct a plot similar to that shown in Figure 14.10 (pH versus volume NaOH added).
Repeat the procedure in Exercise 61, but for the titration of 25.0 mL of 0.100 M NH3 (Kb = 1.8 105) with 0.100 M HCl.
Consider the following pH curves for 100.0 mL of two different acids with the same initial concentration each titrated by 0.10 M NaOH. a. Which plot represents a pH curve of a weak acid, and which plot is for a strong acid? How can you tell? Cite three differences between the plots that help you decide. b. In both cases the pH is relatively constant before the pH changes greatly. Does this mean that at some point in each titration each solution was a buffered solution? c. True or false? The equivalence point volume for each titration is the same. Explain your answer. d. True or false? The pH at the equivalence point for each titration is the same. Explain your answer.
Which of these buffers involving a weak acid HA has the greater resistance to change in pH? Explain your answer. (i) [HA] =0.100 M = [A] (ii) [HA] = 0.300 M = [A]
Consider the titration of 100.0 mL of 0.200 M acetic acid (Ka = 1.8 105) by 0.100 M KOH. Calculate the pH of the resulting solution after the following volumes of KOH have been added. a. 0.0 mL b. 50.0 mL c. 100.0 mL d. 150.0 mL e. 200.0 mL f. 250.0 mL
Determine the volume (in mL) of 1.00 M NaOH that must be added to 250 mL of 0.50 M CH3CO2H to produce a buffer with a pH of 4.50.
A good buffer generally contains relatively equal concentrations of weak acid and conjugate base. If you wanted to buffer a solution at pH = 4.00 or pH = 10.00, how would you decide which weak acidconjugate base or weak baseconjugate acid pair to use? The second characteristic of a good buffer is good buffering capacity. What is the capacity of a buffer? How do the following buffers differ in capacity? How do they differ in pH? 0.01 M acetic acid/0.01 M sodium acetate 0.1 M acetic acid/0.1 M sodium acetate 1.0 M acetic acid/1.0 M sodium acetate
A buffer is composed of formic acid and its conjugate base, the formate ion. (a) What is the pH of a solution that has a formic acid concentration of 0.050 M and a sodium formate concentration of 0.035 M? (b) What must the ratio of acid to conjugate base be to have a pH value 0.50 units higher than the value calculated in part (a)?
A mixture contains initially benzoic acid C6H5CO2(weak acid), Ka = 6.28 x 10^-5, and its concentration is 0.250 M, and sodium benzoate of 0.400M. 500mL of this buffer were titrated with a strong base 250 mL of NaOH that has a pH of 12.875. a) What should the pH be after the addition of the base? b) Why do you think that this pH is correct? Please explain in terms of what really happened during the process that led to this pH, what specific ions or molecules affected the result? c) Sketch the titration curve that will result from this process. d) Predict the pH at the equivalence point, less, more, or equal to 7? Explain?
A 1.0L buffer is made at 25°C that consists of 0.80 mol of a weak acid HA and 0.80 mol of its conjugate base A-. The acid ionizarion constant is Ka= 1×10^-4. To this buffer solution 0.80 mol NaOH is added. Assuming that the volume change due to the added NaOH is negligible, calculate pOH in the resulting solution.
You make a 200mM potassium phosphate buffer (pKa of 7.20) with monobasic potassium phosphate (molecular weight 136.09 g/mol) and dibasic potassium phosphate (174.20 g/mol). For 100mL of buffer, at a pH of 7.20, how many grams of monobasic potassium phosphate are needed? Please walk through this step-by-step.
A 1.0 L buffer solution initially contains 0.25 mol NH3 (Kb=1.8x10^-5) and 0.25 mol of NH4Cl. In order to adjust the pH of the buffer to 8.75 should you add NaOH (MM=40) or HCL (MM=36.5) to buffer the mixture? What mass of the correct reagent should you add to 3 sig figs