The purpose of this lab is to determine the Ka of an indicator within an unknown solution. This can be determined because acid-base indicators are weak bases or acids, and when neutralized they change colors. The color change is measured by the light absorption spectra in the compound which can be used to calculate the concentration. This experiment measures the concentrations by the different pH values which allows the pKa and Ka to be determined. Through the experiment, and the recorded color change with the use of the Spec 20 the unknown Ka of the solution was found.
Introduction
As stated previously because acid-base indicators are weak bases or acids, they change colors. This color change will be used to determine the concentration because …show more content…
To begin the experiment, the buret was prepared for a 0.5 M NaOH and a pH meter was standardized at values 4.00,7.00, and 10.00. Next the 60 ml of 0.1 M HCl was prepared by diluting the 1.0 M HCl provided. Then 5.00 ml of 0.20 g Bromothymol Blue was added as well as a buffer; KH2PO4, 15 ml of H20, and 10 ml 0.1 M HCl which were all measured by a graduated cylinder into each bottle. The pH was then measured and recorded in bottle number one. The next two bottles were buffer solutions and the pH electrode was placed in bottle number two while 0.5 M NaOH was added from the buret while mixing until the reading stabilized to a target pH range. This was then done to all bottles. The spec 20 was then used to measure the absorptions of the four bottles. From the absorption, calculations were able to be made to determine the Ka of the …show more content…
The pKa was determined to be 7.21 while the average Ka was 6.17*10^-8. The Ka % error was -21.91 which means that the yield was lower than expected compared to the literature value which was given to be 7.9*10^-8. (See equation 4)
Results and Discussion In this lab the experimental pKa value was determined to be 7.21 and the average Ka was determined to be 6.17x10^-8. This value shows that through he dissociation of the hydronium in the solution it made the solution more acidic than a basic solution. The previously determined Ka gave a percent error of -21.91%, this error could be due to several reasons including, but limited too: calibrating the ph electrode correctly or possibly spilling the solutions when transferring from bottle to bottle. A negative percent error means that the actual yield was higher than the theoretical yield, which was given at a literature value of 7.9*10^-8.
Comments and
Introduction
As stated previously because acid-base indicators are weak bases or acids, they change colors. This color change will be used to determine the concentration because …show more content…
To begin the experiment, the buret was prepared for a 0.5 M NaOH and a pH meter was standardized at values 4.00,7.00, and 10.00. Next the 60 ml of 0.1 M HCl was prepared by diluting the 1.0 M HCl provided. Then 5.00 ml of 0.20 g Bromothymol Blue was added as well as a buffer; KH2PO4, 15 ml of H20, and 10 ml 0.1 M HCl which were all measured by a graduated cylinder into each bottle. The pH was then measured and recorded in bottle number one. The next two bottles were buffer solutions and the pH electrode was placed in bottle number two while 0.5 M NaOH was added from the buret while mixing until the reading stabilized to a target pH range. This was then done to all bottles. The spec 20 was then used to measure the absorptions of the four bottles. From the absorption, calculations were able to be made to determine the Ka of the …show more content…
The pKa was determined to be 7.21 while the average Ka was 6.17*10^-8. The Ka % error was -21.91 which means that the yield was lower than expected compared to the literature value which was given to be 7.9*10^-8. (See equation 4)
Results and Discussion In this lab the experimental pKa value was determined to be 7.21 and the average Ka was determined to be 6.17x10^-8. This value shows that through he dissociation of the hydronium in the solution it made the solution more acidic than a basic solution. The previously determined Ka gave a percent error of -21.91%, this error could be due to several reasons including, but limited too: calibrating the ph electrode correctly or possibly spilling the solutions when transferring from bottle to bottle. A negative percent error means that the actual yield was higher than the theoretical yield, which was given at a literature value of 7.9*10^-8.
Comments and