7. Repeat all solutions twice for accuracy of results. When cleaning conical flasks, use pure water and remove any precipitate on the glassware with a paper towel.
This experiment has to be fault proof as possible so that accurate results are achieved. To do this, all variables have to remain constant except for the one being measured.
Variables include:
A) Concentration of solution. This will be altered throughout the experiment, but the solution must be accurate to the concentration it represents. To keep it so, burettes are use in measuring out the solutions of water and thiosulphate and volumetric pipettes used for measuring out acid. Both of these pieces of apparatus are accurate to fractions of cm3 and will provide accurate measurements as long as they are taken from the bottom of the meniscus each time.
B) Depth of solution with conical flask. Having different depths of liquid within the flask will create different levels of obscurity and hence affect the times taken for the cross to be obscured in. Hence, we will use the same shaped conical flasks and volumes of solutions each time (55 cm3).
C) The size and shape of the cross has to be the same for each experiment and easily copied so that if the experiment were to be repeated, the same conditions could be applied. Hence, a cross printed from a computer is appropriate. The crosses used were in font Verdana at font size 144.
When results are recorded and graphs plotted from the data, the graphs can be used to calculate the order of the reactions with respect to either reactant. This is either done by plotting concentration against time and taking half-life’s or by plotting rate against concentration and using the shape of the graph to determine the order of the reaction with respect to the reactant measured.
1(page 1, line7) Ratcliff, Brian. Chemistry 2. 2002. Page 109. (Second Order Reactions Section).
Using the method noted above, the following results were obtained:
Where concentration of sodium thiosulphate was altered:
Where concentration of hydrochloric acid was altered:
However, to find the order of the reaction with respect to each reactant, the rate must be plotted against the concentration graphically.
The rate is calculated by taking tangents at each of the points marked on the graphs and measuring their gradients. This is done by dividing the change in y by the change in x [Rate = y/x].
Using µ values plotted from above results for calculating rate:
Looking at the graphs produced from these sets of data, shows us 2 different patterns. The first graph plotted from the thiosulphate reactions, gives us a good indication to a straight line, hence the half-life of the reaction is 2independent of concentration. It shows us that as the concentration is doubled, the rate doubles also.
This is a very good indication as to a first order reaction. Hence this shows that the order of the reaction with respect to sodium thiosulphate is first.
The second of the 2 graphs produced with rate plotted against concentration is from the hydrochloric acid reactions. This line is somewhat more ambiguous mainly because there are some largely anomalous results involved. However, when considering the graph for these set of reactions with concentration plotted against time, it represents straight line which shows that by increasing the concentration it has no effect whatsoever on the rate of the reaction. The order of the reaction with respect to hydrochloric acid is zero.
Using these 2 pieces of information, we can write the rate equation thus:
Rate = k[Na2S2O3][HCl]0
There is still 1 unknown, that is the rate constant k. We need to substitute values of concentration and rate, to calculate k.
Hence,
We shall use the reaction involving 0.020 mol.dm-3 sodium thiosulphate and 0.010 mol.dm-3 hydrochloric acid. The rate for this reaction is 1.30 x 10-3 mol.dm-3.s-1.
0.0013 = k x (0.020) x (0.010)0
k = 0.0013 / 0.020
k = 0.065 = 6.5 x 10-2.
We also need to calculate the units for the rate constant k.
Rearranging:
k = mol.dm-3.s-1 / mol.dm-3 x mol.dm-3
k = mol-1.dm3.s-1.
2(line 3, page 5) Ratcliff, Brian. Chemistry 2. 2002. Page 110. (Half-life and Reaction Rates Section).
Considering the conclusions made from the extrapolated data had to be interpreted; shows that the patterns were not a textbook match. There were some large anomalies in the data as a result of both method itself and techniques used in extrapolating data.
Looking at the first set of results in graph 3, it seems that the anomalous result is with the value for the 0.004 M solution. However, if this was shifted into line with the other results it would give a rate of 0 which considering it has a part in the reaction would not make any sense whatsoever.
The truth of the matter is that the middle 3 results from the data are wrongly distributed. The line of best fit is approximately represented on graph 3.
The reason for these inconsistencies is that the end-point of the reactions are somewhat ambiguous and hence errors can be made in the timing of the reactions. In these cases the change in clarity of the solutions is so great that in some cases the end-point is suspected sooner than it actually occurs. This affects the middle 3 results more than the first and last. The first isn’t affected as much because the time between the suspected end-point and the real one is a relatively small amount of time and will affect the gradient less than with a greater lapse of time. The last result may have falsely been recorded also, but because when dealing with taking tangents to such shallow angles, it makes little difference overall.
On graph 5, we can see a ‘jumble’ of results which bears almost no resemblance to a zero order reaction when rate is plotted against concentration. However, in this case we have to consider its closely related graph; graph 4. This graph is typical of a zero order reaction showing a straight line which shows that no considerable change has occurred when the concentration of the reactant has been increased by up to a factor of 400%. The changes are at most up to 5 seconds in length and when considering a time interval of approximately 38 seconds, this clearly shows a zero order curve.
Graph 5 has a curved appearance due to the fact of the very slight deviations in time on graph 4 which creates a curve in the line and hence tangents about that point take altered gradients.
I believe that graphs 2,3 and 4 show accurate graphical representations of theoretical models. However, graphs 1 and 5 do not show such accuracy. These are to do with both the method of producing results and human techniques in extrapolating data.
Graph 1 shows unequal half-life’s. This and a large amount of the errors in the experiment are due to the method of using the such crude equipment such as a cross disappearing by a precipitate forming. Firstly, the method of using an object being obscured by a solid product is not without its flaws since this is meant to represent the end-point of the reaction and is heavily dependant upon the depth of the solution in the first place. Therefore each time a more dilute solution is used, the time between the calculated and actual end-point grows larger, producing less accurate results.
This is admittedly lessened when a deeper amount of solution is used, however this is at the cost of producing a false end-point. The result of this is that the results are closer to a theoretical model in shape but not in magnitude.
Graph 5 is heavily dependant upon the calculation of tangents at different point on graph 4. The reason why graph 5 is less accurate than its counterpart graph 3, is that the tangents it is measured from change less than in graph 3 and hence are more difficult to draw accurately. This is a large flaw in the method in that it is not scientifically based and neither computer nor calculator can draw the tangents.
I believe graph 5 to be mostly a product of this error.
The main sources of error in this experiment are that is based too heavily upon the ability to accurately record the time interval of a reaction to go to completion by eye. It is inaccurate in that the end-point of the reaction is false represented depending upon the shape of conical flask used and the depth of the solution used.
Compared with using burettes and volumetric pipettes for measuring out correct concentration of solutions, human error is a larger source since the glassware used in a laboratory is made to very high tolerances and human sight is responsible for judging when an object is obscured which when dealing with slower reactions can result in hugely different results each second lowers the overall rate considerably.
The method of timing the intervals was done via a clock which did not work in fractions a seconds and therefore all readings are approximated to integer values. This is an important factor considering the zero order graph being skewed due to it being out by fractions of seconds.
If the experiment were to be repeated, more accurate timing devices would have to be used and co-ordinated accurately with the start of the reaction. Using colorimetry instead of sight, the clarity of the solution could be measured electronically and a stopwatch used to time the change on the meter. This method would remove the reliance on sight and less potential for human error which would otherwise prove problematic.
Otherwise, if the experiment were to be repeated, the method would have to be completely altered in order to reduce error.
Firstly, the concentrations would have to be lowered in order to produce a solution which would form the precipitate slower and therefore the change could be measured more accurately.
I would prepare solutions of the same concentration and quench the reactions at different time intervals in order to deduce the concentration change with respect to time. 3Using an iodine solution (e.g. iodine in potassium iodide) would mean that I- ions would compete with the H+ ions (acid) for thiosulphate ions and immediately quench the reaction. Using an equal molar quantity with what started in the reacting flask of iodine, and then staining the quenched reaction with starch would produce a blue/black colour which when using colorimetry could be used to calculate amount of iodine left in solution and hence proportional to the amount of thiosulphate that has reacted.
This method of carrying out the experiment means that there is little possibility of creating anomalous results so long as the concentrations of the solutions are measured out accurately using accurate equipment like the burette and volumetric pipette.
3(page7, line 37) <http://members.rogers.com/acidmanual2/analytical_so2acid.htm>
