The determination of the rate equation between

The determination of the rate equation between
the reaction of HCl and Na2S2O3

Aim

To determine the rate equation between the reaction of known concentrations of HCl and Na2S2O3.

Scientific Background

To carry out this task we need to find the initial rate of reactions of different concentrations of the substances. We are supplied with the following chemicals, all of known concentration:

0.4 mol dm-3 sodium thiosulphate (Na2S2O3)

2.0 mol dm-3 hydrochloric acid (HCl)

Deionised water (H2O)

2HCl(aq) + Na2S2O3(aq) → 2NaCl(aq) + SO2(g) + S(s) + H2O(l)

Many conditions affect the initial rates of reaction including the concentrations of the substances, temperature, total volume and pressure for example. This is all described in collision theory where for molecules to react, they must collide and have the required energy for the reaction. All of the factors will affect the probability of the molecules colliding.

Concentration: this is effectively the amount of the molecules in a certain volume. If there is a low concentration, and therefore are few of the molecules in a certain volume, there will be few of them to collide so the rate of reaction will differ due to its concentration.

Temperature: this determines how fast the molecules are moving. If they are moving faster then they will collide more which will increase the chances of them making a successful collision and reaction. It is also the energy the molecule has and if its energy is enough then it will be able to react otherwise there will be no reaction.
In a mixture of molecules, there may be some which do not have the required energy to react and some which do. The proportion of these molecules can be shown by the Maxwell-Boltzmann distribution curve. As shown on the graph, all the molecules in the shaded area have the required energy (activation energy) for the reaction and will react if it makes a successful collision.

http://www.webchem.net/notes/how_far/kinetics/maxwel2.gif

The rate equation for this reaction will be in the form of:

Rate = Kc x [HCl]x x [Na2S2O3]y x [H2O]z

It will be in this form because all three of these chemicals are present in the mixture during the reaction. The orders of these chemicals may be 0 meaning they are not in the slow step of the mechanism, or they may have an order above 0 where they will be present in the slow step. The order of that chemical determines how many molecules are in that stage.

Catalysts are molecules that can reduce the activation energy required for certain other molecules. The presence of a catalyst will speed up the rate of reaction due to the larger amount of molecules that have the required activation energy so can react. These catalysts can be take part in the reaction itself however the catalyst is always regenerated so is not part of the general equation. Autocatalysis is when catalysts for a reaction are produced by that reaction. This will cause the rate of reaction to increase as more and more of the catalyst is produced and this rate of production will also get faster and faster. This is a form of positive feedback and the rate will keep increasing until a limiting factor occurs and the rate levels off. Catalysts can not affect the amount of reactants produced.

Variables

Only the concentrations of the solutions should be varied so that the rate equation can be determined.

Concentrations of Solutions (controlled variable): These will all be varied and compared to each other so that their individual orders can be determined. By keeping them all constant and varying only one at a time this can be achieved. The concentration of water will not be measured and will be assumed to be constant.

Time for Reaction (dependant variable): This will be the time until the reaction has caused the marking below the conical flask to be obscured completely so it is no longer visible.

Volume of Solutions (constant): This will be kept constant to keep it a fair test. If this did vary then the depth of the solution would vary in the conical flasks and the points at which the marking underneath the flask will become obscured would vary making it an unfair test. It will be kept at 75cm3.

Temperature (constant): This affects the rate of reaction so must be kept ...

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Time for Reaction (dependant variable): This will be the time until the reaction has caused the marking below the conical flask to be obscured completely so it is no longer visible.

Temperature (constant): This affects the rate of reaction so must be kept a constant. It will be kept at room temperature.

Hypothesis

I believe that the higher the concentration of each solution, the higher the rate of reaction due to a larger amount of the reacting molecules and therefore the larger amount of successful collisions made by them. Increasing both concentrations individually will increase the rate of reaction. I think that HCl will be to the second power due to the larger amount of this molecule and that Na2S2O3 will only be first power.

Prediction

The higher the concentration of either solution will increase the rate of reaction. As the concentration of Na2S2O3 doubles, the rate will double. As the concentration of HCl doubles, the rate will quadruple.

Safety

Before any of the experiments are carried out, these safety points must be read. Must wear safety goggles when near any chemicals.

HCl is in a high concentration so gloves must be worn when using it.

Keep the work area safe and tidy; when equipment is no longer required, move it aside or clear it away.

To avoid spillage; ensure the burette tap is closed before filling and use a funnel. Remove the funnel from the burette when not in use, this could cause a hazard and any chemicals dripping off the funnel into the burette will cause inaccurate results. Wash out all equipment before and after use to remove all traces of chemicals which could cause errors in the experiment.

Preliminary

a) Aim
The aim of the preliminary experiment is to find a suitable range of concentrations to use which will give a range of times which are not too close together and not too far apart so that the results can be processed correctly.

b) Apparatus

c) Method

1. Clean all equipment with de-ionised water to avoid contamination which could lead to inaccurate results. Fill all of the burettes with their designated solution to their 0 mark.

2. Measure out an amount of HCl and Na2S2O3 from the burettes into the 100ml beakers. The volumes of these two solutions added together must not exceed 75cm3.

3. Measure out an amount of deionised water in its own 100ml beaker to make the total volumes of all of the solutions added together 75cm3.

4. Add the water and the HCl to the conical flask and shake gently. Add the Na2S2O3 to the conical flask, start the timer and shake gently then put the conical flask over the tile with the mark on.

5. Look down through the conical flask at the mark and as soon as the mark is no longer visible, stop the timer. Write down the volumes of the solutions and the time for processing.

6. Repeat steps 1 to 5 until a large set of different volumes have been used.

d) Diagram

e) Results and Processing

To find the rate, it is simply 1/Time. Concentrations of the solutions can be calculated by simple formulae from their dilutions. To find the amount of moles of the chemical it is Initial Concentration multiplied by the volume of it collected from the burette. All units have to be standard (i.e. in dm3, 1dm3 = 1000cm3). The initial concentrations are 0.4 moldm-3 and 2.0 moldm-3 for Na2S2O3 and HCl respectfully. To then find the concentration it is Amount of moles divided by Total volume which is 75cm3. Final Concentration = Initial Concentration x Initial Volume / 0.075dm3.

E.g. for the first value of HCl on the table:
Final Concentration = (2.0 x (2.5/1000)) / 0.075 = 0.005 / 0.075 = 0.067 moldm-3

Preliminary Analysis

From this set of results I have determined a suitable range of volumes of the solutions. 2.50cm3 of HCl is too low but 5.00cm3 and above should be suitable for HCl. For Na2S2O3 50.00cm3 is not too fast and its lower values are not too slow so a large range can be taken for this solution.

From these results we can see that as the concentration of Na2S2O3 doubles, the rate doubles which would suggest it is to the first order. However as the concentration of HCl increases there is only a slight increase to the rate which is not even close to doubling or quadrupling as predicted.

A major problem did arise from the method, when adding the Na2S2O3 to the conical flask it was difficult to start the timer at exactly the same time and there was a slight delay for to timer to be engaged. A solution to this would be to get help from somebody else to start the timer for you.

Final Apparatus

Final Method

1. Clean all equipment with de-ionised water to avoid contamination which could lead to inaccurate results. Fill all of the burettes with their designated solution to their 0 mark.

2. Measure out an amount of HCl and Na2S2O3 from the burettes into the 100ml beakers. The volumes of these two solutions added together must not exceed 75cm3.

3. Measure out an amount of deionised water in its own 100ml beaker to make the total volumes of all of the solutions added together 75cm3.

4. Add the water and the HCl to the conical flask and shake gently. Add the Na2S2O3 to the conical flask, get an assistant to start the timer and then shake the flask gently and put it on the tile with the mark on.

5. Look down through the conical flask at the mark and as soon as the mark is no longer visible, stop the timer. Write down the volumes of the solutions and the time for processing.

6. Repeat steps 1 to 5 until a large range of results have been collected. Take three repeats for each set of volumes of the solutions. For one set take increments of 10 from 10 to 50cm3 for Na2S2O3 with 5cm3 of HCl and for another set take increments of 5 from 5 to 25cm3 for HCl with 50cm3 of Na2S2O3.

Final Diagram

Results

Analysis

The table displays all of the processed results showing the rates and concentrations of the different mixtures. The two sets of results were taken to see what effect each different chemical has on the rate of reaction which is displayed on the graphs.

The graph showing the rate against HCl is not normal due to the autocatalysis which is taking place. It is very steep to begin with however it levels off showing that it may be 0 order. The Na2S2O3 graph shows that this is defiantly 1st order so it must be the only thing present in the first reaction as well as the catalyst which is regenerated so a set of reactions can be suggested.

2HCl(aq) → 2H+(aq) + 2Cl-(aq) (HCl dissociates in water)

Na2S2O3(aq) + 2H+(aq) → H2S2O3(aq) + 2Na+(aq) (slow step)

2Na+(aq) + 2Cl-(aq) → 2NaCl(aq) (fast step)

H2S2O3(aq) → H2O(l) + S(s) + SO2(g) (fast step)

SO2(g) + H2O(l) → H2S03(aq) → HSO3-(aq) + H+(aq) (H+ formed, autocatalysis)

H+ must be the catalyst in this case and there is a possibility that this set of reactions could be how the chemicals actually react. The order of the Na2S2O3 is as I predicted however the order of the HCl is not, however the orders I have suggested may not necessarily be correct.

Evaluation

There are anomalies on the tables which are highlighted and one anomaly on the graph of Na2S2O3 against Rate which is also highlighted. These anomalies have been excluded from the averages taken during the processing of the results. They could have occurred for many different reasons but it is most likely that they came from not starting the timer at the same time as the reaction started which could have put these results out by as much as a few seconds. This probably was not the reason for the anomaly found for the rate, as shown on the table and graph, as the repeat readings should have eliminated this problem and the fact that the values for time should be 220 seconds to fit the trend line which is very different from the results observed.

As previously mentioned, there is a major problem in starting the timer at the exact moment. It is difficult to start the timer while mixing the chemicals yourself so assistance must be acquired. The results may also be affected by residue left on the surface of the beakers and the flask. This may contaminate other solutions if the beakers are not washed out thoroughly however, apart from these minor problems, there are very few improvements that could be made to the equipment apart from using a full data-logging system. The results can also be affected by ambient temperature due to temperature and pressure having an effect on reaction rate. This can be avoided by using a controlled environment or by ensuring that these conditions are the same by measuring them before the experiment takes place. Doing everything on the same day will make it easier to keep these conditions the same.

As in all experiments it would be better to take many more repeat readings and have a larger range of readings to see the full extent of the concentration vs. rate graphs. I believe that this set of results was adequate to come to the conclusion of the orders of the reactants however the HCl vs. Rate graph was overcomplicated with the autocatalysis affecting it and may have been interpreted incorrectly.

Rate = Kc[Na2S2O3]1[HCl]0[H2O]0

Bibliography

Maxwell-Boltzmann distribution curve

http://www.webchem.net/notes/how_far/kinetics/maxwel2.gif

Reaction Orders

Advanced Sciences – Chemistry 2 – Brian Ratcliff, Helen Eccles.

Catalysts

Advanced Sciences – Chemistry 1 – Brian Ratcliff, Helen Eccles, David Johnson,
John Nicholson, John Raffan

Autocatalysis

http://www.fact-index.com/a/au/autocatalysis.html

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