Lab Report 1

Lab 1: Buffers Introduction:

The central purpose of Lab 1 was to demonstrate the power and importance of buffers in regulating the pH of solutions. By adding very small amounts of acid or base to different  buffer solutions, students were able to calculate the buffer buffer capacities of those solutions and some general trends with regard to buffer strength with regards to acid/base balance. The students also learned useful skills applicable to laboratory settings in general. One relationship in particular, particular, the buffer capacity β, was of central importance in evaluating the  buffers used in the lab. The buffer capacity was calculated using the equation:

  

d [b] 

dpH 

d [a] 

dpH 

In this equation β represents represents the buffer capacity, d[b] and d[a] represent the very small change in the concentration of a strong base or acid respectively, and dpH represents the very small change in pH resulting from the addition of strong acid or base. Understanding  buffer capacity is important because it allows allows chemists to determine wether or not a given  buffer is suitable suitable for an application and what the limits of that buffer buffer are.

Procedure:

Dissolve 1.742g K 2HPO4 and 1.361g KH2PO4 in separate 100mL volumetric flasks to  produce 100mL of 0.1 M solution of each. Empt y the flasks into separate, labeled 250mL  beakers. To make solutions 1-5 mix the K 2HPO4  and KH2PO4 solutions in 40:0, 30:10, 20:20: 10:30; and 0:40 mL ratios and dilute all with 60mL distilled water. Also prepare 100mL distilled water as solution six. Mix each solution before taking and recording its

initial pH. Pour 20 mL of each solution into a separate beaker. Add 0.25 mL of HCl to each solution. The concentration of HCl needed to affect a small but measurable change in pH varies between solution. 0.05 M HCl should be used for solutions 1 and 5 and 0.5 M HCl should be used for solution 3. Solutions 2 and 4 should take intermediate concentrations around 0.1 M. Add very dilute HCl (around 0.01 M HCl) to the water solution. If the pH change it too lager or too small (0.1 < dpH < 0.5) then pour a new 20mL sample and use an appropriately adjusted concentration of HCl. Stir the solution and record the pH after the addition of the HCl. Repeat the procedure of the paragraph above but use a strong base, NaOH, instead of the HCl. Add ~0.2g of aspirin and Bufferin to separate, labeled 150mL beakers and attempt to dissolve the aspirin and Bufferin, each in 25 mL water. Record the pH of the two solutions. Add 0.25mL 0.5M HCl to each beaker and record the pH. Make two new beakers and aspirin and Bufferin solution and add the same amount of NaOH. Record the pH. Finally,  properly dispose of all waste.

Results:

See Appendix 1 for all tables and graphs.

Discussion:

Though the group only had time to calculate a very small amount of data, the data

collected was sufficient for the central focus of the experiment. The experimental data confirms what the theoretical data predicted: that the buffering capacity of the solution would be greatest at roughly equivalent concentrations of weak acid and conjugate base. Experimental data shows that the buffer capacity β of solution 3 was roughly ten times that of solutions 1 and 5, with solution 3 having the closest balance of acid and conjugate base. The buffer capacity was effectively measured to a useful degree for predicting future changes. The data collected answers the central questions posed by the lab and the  phenomena discussed in the pre lab can all be seen through the collected data. The graph of buffer capacity across a pH range from 4 to 9 shows that the buffer system is most effective around pH 7, though the theoretical values would put the peak around 7.3. If HCl were used to protonate KH2PO4 to KH3PO4 a new buffer system could  be made which would be effective at a lower pH. Though the general trends in the collected data were expected the measured pH values varied greatly from the theoretical values, at times more than one pH. This variation could have been a result of mis-calibration of pH probes, however it seemed that many others had similar discrepancies. The error may also have arisen from impurities in the solutes dissolved in the solution. The group had a great deal of difficulty performing the experiment in a timely manner and was not able to perform every test on every solution. The data collected has a wide  breadth and is reliable. It shows trends well, however its predictive ability is lessened by the lack of a larger data set. Conclusion:

Despite difficulties in data collection, the group found that the buffer capacity of the

most balanced solution was significantly larger than the two others tested. The concept of  buffer capacity was effectively used in the experiment. It was found that despite sometimes significant numerical variation from predicted values the solutions behaved largely as  predicted.

APPENDIX 1 Solution

Initial pH

pH after H+ addition

pH after OHaddition

Buffer capacity [H]

Buffer capacity [OH]

1

8.98 8.62 (0.05 M) 9.36 (0.05 M) 0.00174 0.00164 9.91 9.00 (0.05 M) 10.8 (0.05 M) 0.000687  0.000702 7.69 7.55 (0.2 M) 7.85 (0.2 M) 0.0179 0.0156 2 3 6.91 6.56 (0.5 M) 7.26 (0.5 M) 0.0179 0.0179 7.21 6.93 (0.5 M) 7.49 (0.5 M) 0.0223 0.0223 6.73 6.57 (0.2 M) 6.87 (0.2 M) 0.0156 0.0179 4 3.20 (0.05 M) 0.000393 5 4.79 5.22 (0.05 M) 0.00145 4.30 5.41 (0.05 M) 0.000563 7.00 3.20 (0.05 M) 10.80 (0.05 M) 0.000164 0.000164 DI Water There was insufficient data given about aspirin and Bufferin to make useful Aspirin Bufferin  predictions in the absence of any experimental data  Note: values values in italics italics are theoretical theoretical values values

BufferCapacityoftheHP BufferCapac ityoftheHPO4/H2PO4syst O4/H2PO4systematvar ematvar ryingpH 0.025      y        t         i      c      a       p      a         C       r      e         f         f       u         B

0.02 0.015

Theoretical [H+] Experimental

0.01 0.005 0 4

5

6

7 pH

8

9

10