Table showing my personal raw data:
Table showing the raw data of the entire class:
Data processing:
Table showing the average time taken for the neutralization point to occur
Sample calculations :
Average time taken for the neutralization point to occur:
(data of group 1 + group 2 + group 3 + group 4) / (number of groups)
For example, for the average time for 3mL of urease added:
(35 + 40 + 42 + 49) / 4 = 41.5 min = 41 min 30 sec
Conclusion and Evaluation:
The universal indicator can be used in this experiment to measure the rate of the reaction, because urea is broken down according to the following reaction:
Urea ammonia + carbonic acid
Ammonia is very alkaline. The solutions in this lab always contained 2mL of CH3COOH(aq)
(carbonic acid). Hence, a neutralisation reaction takes place as ammonia is produced from the breakdown of urea. When all carbonic acid molecules have reacted with ammonia, the neutralization point is reached and the universal indicator turns from pink (color displayed in acidic conditions) to green (color displayed in basic conditions).
The higher the urease concentration, the closer are the quarters in which the urease molecules are traveling and hence, the more collisions between the substrate molecules (urea) and urease will occur: The rate of the reaction is faster. Hence, the hypothesis is that with increasing concentration of urease, the time it takes to neutralize the ethanoic acid will decrease. Looking at graph 1, the results of this experiment support the hypothesis. The average time taken for the neutralization point to occur decreases constantly with increasing urea concentration. The best fit line is of the form ax2 + bx + c and hence, the time first decreases very fast and then gradually slows down with increasing urease concentration, that is, the graph flattens. This means that the rate of the reaction increases fast at first and then slows down. It slows down because at some point all substrate molecules will almost immediately have bound to an urease’s active site as soon as the urease is added to the urea and hence, a further increase in urease concentration will not affect the rate of the reaction any more. Hence, one can say that the results do reflect pretty well the theory and can therefore be considered fairly reliable.
Although the expected results were in fact obtained in this experiment, the procedure is certainly not perfect and contains experimental errors. First of all, we have the uncertainties caused by the apparatus, more specifically by the measuring cylinder used. The uncertainty of the watch has been ignored here, because its uncertainty (±1 second) is negligible here where the value for the time is 19 minutes and more.
Furthermore, there is the error involved in the fact of using an indicator to determine the neutralization point. The point at which the colour change occurs is actually already beyond the point of neutralization. In the case of this lab, when the universal indicator turns from pink to green, the solution has already a pH higher than 7 and not of exactly 7. This causes the values for the time taken to reach the neutralization point to be higher than they should have been. However, this error is involved in all the reactions of this experiment equally, hence, when comparing these values for solutions with different concentrations of urease, the error caused by this is much less significant.
Moreover, the colorchange could not be detected immediately, because although the test tubes were shaked constantly, there is a certain delay before the colorchange is visible in the entire test tube. This causes the values for the time taken to reach the neutralization point to be slightly more than it should be. A slight improvement in the procedure would probably be to use a magnetic stirrer instead of shaking the solution by hand. It would reduce a little this delay in time before the colorchange can be detected.
Very obviously, there are certain errors involved in using an indicator as a method to measure the rate of the reaction. An improvement of this experiment would be to measure the pH of the solution, using a pH metre, at different time intervals, for example, every 10 seconds. This would have made it possible to plot a graph of time versus pH for each solution where the slope of the graph would represent the rate of the reaction. In this way, a much more thorough comparison could have been drawn between the rate of reaction and urease concentration. Apart from comparing at what time the solution reaches pH 7, one could for example also compare the value of the maximum instantaneous rate (where the slope of the graph is the highest) and after what time it occurs. The fact of using a pH metre to measure the rate of reaction would also eliminate the error of the colorchange occuring actually after the neutralization point and not right at the neutralization point. The simplest improvement would be to perform more trials, which would reduce the experimental error in the average value.