By: Howard Andres

In this experiment, our lab group had several purposes to achieve.  One purpose of the lab was to determine the approximate pH of several solutions by using acid base indicators.  Then our group had to determine the pH of some other solutions with a pH meter.  And finally, the third purpose of our lab was to note the effect on pH by changing the composition of a buffer, such as adding small amounts of acid and base to the buffer, diluting the buffer, and destroying the buffer.  From the data we would obtain, our group would then be able to determine the Ka of the acid.

In the first part of our experiment our group was to use the observations we obtained through using different acid base indicators on several different solutions, which included HCl, NaH₂PO₄, HC₂H₃O₂, and ZnSO₄ to determine their approximate pH.  With HCl, the colors produced by the indicators of congo red and bromcresol green provided the best possible pH range of about 3-4.  In NaH₂PO₄, thymol blue and congo red provided the most suitable pH range of 4-5.  The colors produced by methyl violet and thymol blue provided the best pH range for  HC₂H₃O₂ at about 2-3, and congo red and bromcresol green provided a good pH range of about 5-6 for ZnSO₄.  An interesting observation to note for this part of the lab is that thymol blue produced a light blue color with HCl, a color that does not seem to be in the range of thymol blue’s red to yellow.  Some kind of possible error is hinted here because it, a strong acid, had a pH range higher than that of the weak acid HC₂H₃O₂.  Either something was wrong with the indicators or our group performed some other mistake to reach such a result.

In the second part of our experiment our group used pH meters to determine the pH of NaCl, Na₂CO₃, NaC₂H₃O₂, and NaHSO₄.  Additionally, we added some of the indicator bromcresol green to help us in making sure our results were satisfactory.  NaCl produced a pH of about 7.68, Na₂CO₃ a pH of about 10.54, NaC₂H₃O₂ a pH of about 7.71, and NaHSO₄ a pH of about 1.54.  Even though these results were not in the range of bromcresol green’s 3-6 pH range, the indicator changed to its most extreme colors so to prove these results were about right, turning yellow for NaHSO₄ since its pH was below 3, and blue for all the other solutions which were above the pH of 6.

In the third part of the lab our group experimented with the buffer system of HC₂H₃O₂, acetic acid, and its conjugate base C₂H₃O₂^-.  By mixing the same amounts of each component (.01M and 15 mL each), the pH of the system was found to be at around 4.60.  Through calculation, [H⁺] was 2.51x 10^-5M and the Ka was 2.51x 10^-5.  When the buffer system was diluted with 30 mL of water, the pH was 4.52, with a [H⁺] of 3.02x 10^-5M and Ka of 3.02x 10^-5.   There seems to be some error with this though, because diluting a solution should not change a buffer’s pH since the concentrations of acid and conjugate base would still be the same.  Most likely, such an error resulted from the tap water used to dilute the buffer, which is probably slightly acidic.  Distilled water was not used for the reason that the instructions did not specifically say to do so.  Then, when 5 drops of NaOH was added the pH went up to 4.57 from the pH of 4.52, and down to 4.51 when 5 drops of HCl was added, a reasonable pH since the addition of the acid brought the pH back down to near its starting point of 4.52.  When [HC₂H₃O₂]/[C₂H₃O₂^-] equaled .10, the pH was found to be at 5.70, and the Ka was 2.00x 10^-5.  When the buffer was broken when excess NaOH was added, the pH went very high, as expected, to 11.23.

When 25mL of distilled water was put into a beaker, the pH was found to be at 4.70 with the pH meter.  Most likely, there is some sort of error with the meter, because distilled water should be at a pH of 7.  It can be concluded then that most likely the rest of the results obtained with the pH meter would be off.  The pH of the water after 5 drops of HCl was 3.05, and 9.57 after 10 drops of NaOH was added.

[ Back ] [ Up ] [ Next ]