Chemical reactions take this analogy even further. Supposing the rate of step one in a particular reaction is of first order with respect to the concentration of chemical A, but the rate of step two is of zeroth order with respect to chemical A, the overall rate will still be of zeroth order, because the intermediate product formed (input to step two) can only be processed at a certain fixed rate (hence it is of zeroth order with respect to chemical A). Any excess intermediate product will remain in the intermediate stage or if it is unstable will revert back to the original chemical until it can be processed in stage two. When predicting the overall rate of reaction, always take the slowest rate from of all steps of the reaction and lowest order with respect to the concentration of a particular chemical.
Endothermic and Exothermic reactions. In normal circumstances reactions happen because it is better for the reaction to happen which allows the molecules would then be put under a more stable state by possessing less chemical energy (Energy stored in bonds, lattices, etc.) The only case where a reaction can still happen despite having a positive reaction enthalpy is when the products have a much greater entropy than the reactants and the reaction progresses endothermically. This is because the world is constantly seeking to increase the entropy, or the messy-ness of the molecules. The more different ways the molecules can be arranged in that state, the greater entropy it has. This reaction is exothermic, but as the rise in temperature is so small its effects can be neglected.
The Activation Enthalpy. We have continually referred to the activation enthalpy. Why should there be an activation energy? This is, in fact, the energy needed to break the necessary bonds upon start of the reaction to enable the reaction to progress. At a particular temperature the energies of each individual molecules are distributed according to the Boltzmann's distribution, where if you plotted number of molecules against energy, there will be a peak representing the average energy (which is proportional to temperature) of the molecules but also molecules with either more or less energy on either side (graph on next page). There would then, theoretically, be a finite number of molecules with more than the activation energy at any given temperature and hence the reaction should occur at any temperature. However, in practise, it is either too slow to be measurable, or a necessary chain-reaction cannot start because of the low rate of the initiating reaction.
To calculate the Activation Enthalpy of a particular reaction, we needed to make use of the Arrhenius Equation. By rearranging Equation 5: ‘k = A exp( -EA/RT )’ into the form of y = mx + c, we are able to relate it to a graph. k represents the ‘rate constant’ which at a given temperature equals to the rate of the reaction, which is directly proportional to t-1, where t is the amount of time taken for the reaction mixture to turn blue in this particular case. If we plot 1/Temperature against log(1/Time) (algebraically equivalent to -log(Time)), the gradient of the graph will be equal to -EA/R. Using this method, we can calculate the activation enthalpy of any particular reaction just by running it a several different temperatures and plotting a graph.
The glad tidings brought by the catalyst. When a catalyst is added it helps the reaction to go faster because it lowers the activation enthalpy. The reactants will bind to the catalyst and this often weakens a vital bond, hence decreasing the energy required to break them so a greater proportion of the reactants can react. In the case of heterogeneous catalysis they also provide a surface on which the reactants can react, and increase the chances of molecules colliding because some molecules can be temporarily 'held' at the surface whilst the other molecule approaches. Enzyme-based catalysts work by allowing the substrate to 'fit' and bind to itself in a strained arrangement, and hence weakening the bond. Most solid catalysts work by providing a surface on which the reactants can react, and other transition mental catalysts work by providing a easier mechanism for the reaction using the catalyst's transition properties.
Predicting the performance of a transition metal catalyst. For the transition metal catalysts that work by acting as an oxidising agent of the substrate then a reducing agent of the product, or vice-versa, whether one particular metal will work can be predicted by using the tables of standard electrode potentials (or oxidation potentials). If the higher state of the catalyst will oxidise the substrate and the lower state of the catalyst can be oxidised by the product (hence reducing the product) then there is a possibility that the catalyst may work. Whether it actually works or not and its efficiency cannot be predicted at this stage, and must be found out by experiment. However, those outside the range of possible catalysts definitely won't work as it requires an input of energy just to carry out the first step, ie to oxidise the substrate.
In this particular instance, the metal catalyst works by utilising its transition properties to cause a chain-reactions which is much faster than the original reaction. In the following example cupric ions (Cu2+) is used as an example, but the same principle can be applied to all transition metal catalysts used in this reaction. The equation for the first step of the catalysed reaction is as follows:
2I-(aq) + 2Cu2+(aq) ---> I2 (aq) + 2Cu+(aq)
2 I-(aq) + 2 Cu2+(aq) > I2 (aq) + 2 Cu+(aq) ........................ Equation 6
The cupric ion is first of all reduced by iodide to cuprous ion, and our end-product iodine is produced in the first step. As we are using starch to detect the presence of iodine, the reaction need not to have completed to its full extent before the presence of the product will be detected. The cuprous ion is then oxidised by the persulphate ion back to the +2 oxidation state, as shown:
2Cu+(aq) + S2O82-(aq) ---> 2Cu2+(aq) + 2SO42-(aq)
2 Cu+(aq) + S2O82-(aq) > 2 Cu2+ (aq) + 2 SO42-(aq) ........................ Equation 7
As both steps involve the collision between a cation and an anion, due to the natural electromagnetic attraction between the two, it happens much often than the single collision between the two anions in an uncatalysed reaction, thus greatly increasing the pre-exponential factor in the Arrhenius equation. However, it should do little to alter the activation enthalpy as this is an ionic reaction and there are no breakage of strong covalent bonds involved. Previous researchA1 indicates that the activation enthalpy of both the catalysed reaction using cupric ions and uncatalysed reaction should be similarA1.
If ferrous ions were used instead, the catalysis would not be as efficient, as the ferrous ions need to be oxidised to the ferric state by persulphate first, and then the ferric ion can then oxidise the iodide to iodine. This reverses the two steps of the reaction, and the iodine will be produced in the second step.
The catalysts to use need to be chosen carefully. It must have a redox potential between the two half-reactions to allow itself to be reduced by iodide and then oxidised by persulphate. By inspecting a redox potential table, I have identified several potential catalysts:
I have decided to try several other catalysts to confirm that they would not display catalytic properties in this reaction. The catalysts identified here are only the ones that could work, whether it actually works or not has to do with the entropy of its other state aswell. For example, solid silver metal has a much lower entropy than its ionic counterparts, and therefore silver ions may not act as a catalyst in this reaction as it would require silver in the solid state be produced in an intermittent stage of the reaction. As systems naturally go towards the side with higher entropy the silver should not allow itself be reduced by iodide ions.
The initial rate of reaction. As in most reactions the rate of reaction is dependent on the concentration of the chemicals it is the initial rate of reaction that we are interested in, because that's when the rate of reaction is a true reflection of the rate given by that particular concentration. This is partly why the iodine-clock reaction and the peroxydisulphate ion is chosen as a sample, as we are able to delay the expiry of the clock by adding Sodium Thiosulphate to the mixture to mop-up any iodine produced and reduce it back to iodide. The equation for this side-reaction is as follows:
I2 (aq) + 2Na2S2O3 (aq) ---> 2Na+I-(aq) + Na2S4O6 (aq)
I2(aq) + 2 Na2S2O3(aq) > 2 Na+I- (aq) + Na2S4O6(aq) ........................ Equation 8
Note that iodide ions (I-) are produced again on the right hand side as Sodium Iodide, replenishing the supply of iodide ions in the reaction mixture. This enables us to maintain an approximate concentration of iodide ions and enables us to measure the “initial rate” for a period, before all the Thiosulphate ions are used up. This technique will only work if the reaction of Thiosulphate with iodine is much faster than the oxidation of iodide by the persulphate ions. This appears to be the case, and may go part of the way to explain why erratic readings are sometimes obtained in reactions with Cupric ions as catalyst, because the catalyst is extremely efficient in this case, and this side-reaction may not be able to keep up with the rate at which iodine is being produced.
Experimental Schedule (for the four weeks)
Week 1 (1) Establish the concentration and optimum range of variables (including maximum available concentration of K2S2O8 available, as K2S2O8 isn’t terribly soluble in water.)
(2) Establish the bailout point (ie the point where I call it ‘gone black’.) It was eventually decided that as soon as I can detect a trace of blue in the reaction mixture I should call it a bailout because that is when the reaction seizes to occur at the initial rate, and waiting for the iodine concentration to build up is not necessarily the best idea.
(3) Plot initial graphs to see if the results are as expected. Get 'Plot' working on Archmedies so that I can just enter the figure and have the activation enthalpy calculated (possibly using PipeDream).
(4) Calculate the standard activation enthalpy at standard concentration, without any catalyst.
(5) Vary the concentration and investigate the effect on the initial rate of reaction.
Weeks 2 & 3 (1) Add Cupric ions catalyst, decide which Copper Salt to use.
(2) Investigate the effect of using different catalysts, including
Ferrous ions (another transition metal)
Stannous ions (a group 4 'metal' with multiple oxidation states)
Mercuric ions (transition metal in +2 oxidation state)
Silver ions (transition metal in +1 oxidation state)
(3) Analyse the results with the different catalysts - and produce alternative hypothesis. As we are expecting the Activation enthalpy to be similar, an alternative explanation of the mechanism of the catalysis should be sought.
Week 4 (1) Use the best catalyst available and investigate the effect of changing the concentration of the catalyst.
Risk Assessment
At 0.2M KI and the other salts used is reasonably safe but precautions associated with dilute salt solution must be taken, ie, the solution should not be ingested or injected. Spillage should be washed with plenty of water. Some salts such as Cupric Sulphate may be harmful to the skin, but at 0.1M as long as there are no prolonged contact the risk associated with the use of such solution is negligible.
Mercury Salts are often poisonous, and solution of Mercury Salts more concentrated than 0.002M should be labelled as so. The long-term effects of Mercury exposure are well documented, and includes mental retardation. However, this experiment does not involve prolonged usage, and the small quantities that I would be using can be washed down the sink with plenty of water when no longer required, or if a leakage occurs. Gloves should be worn to minimise the risk of skin-contact.
K2S2O8 in the solid state represent a health hazard, because it is an irritant. When making up standard solutions gloves and eye-protection should be worn to prevent any possible body contact. If in the unlikely circumstances that it is spilled it should be washed away with plenty of water. Where it is impossible to wash with water (such as near the electronic balance) it should be cleaned using a vacuum cleaner and then wiped thoroughly with a damp cloth.
Silver Nitrate is an irritant and will stain clothing at 0.1M. Protective gloves should be worn when handling the solution and lab coats should always be worn in the lab.
None of the above chemicals should be ingested. If gloves are not used for the dilute solution, hands should be washed thoroughly with soap and warm water at the end of the practical session in order to minimise the risk of ingestion after the lab session.
Materials
(1) Potassium Iodide (KI(aq)) at 0.2M (1 litre)
(2) Sodium Thiosulphate (Na2S2O3(aq)) at 0.01M (500 ml)
(3) Starch Suspension (300 ml)
(4) Potassium Peroxydisulphate (K2S2O8(s)) Powder (100g)
(5) Cupric Sulphate Solution (CuSO4(aq)) at 0.1M (100 ml)
(6) Ferrous Sulphate Solution (FeSO4(aq)) at 0.1M (100 ml)
(7) Stannous Sulphate Solution (SnSO4(aq)) at 0.1M (100 ml)
(8) Mercuric Chloride Solution (HgCl2(aq)) at 0.1M (100 ml)
(9) Silver Nitrate Solution (AgNO3(aq)) at 0.1M (100 ml)
(10) Cobaltous Sulphate (CoSO4(aq)) Solution at 0.1M (100 ml)
Method
Preparation of standard solutions. As there were may people doing different projects in the lab at the same time it was not possible for some of the solution to be supplied. Most of the solutions had to be made up in the lab. The solutions were all in easily dissolvable concentrations therefore it suffices to say that it is advisable to dissolve the solid in a beaker with distilled water before pouring it into the standard bottle as with this method large crystals could be crushed hence it will take less time to dissolve. The beaker must be rinsed out thoroughly and the contents emptied completely into the standard bottle before the standard bottle is topped up with more distilled water.
The starch solution should be kept in the fridge or made up at the beginning of each practical session, because the starch is rather nutritious and forms a good broth for bacteria to thrive on. To make the starch solution 100 ml of water was boiled and then a starch paste consisting of 1g of starch in a little water was poured into it, and then the bottle was cooled under the tap quickly. This makes 1% clear starch solution suitable for indicator purposes, and has the effect of stopping the starch from settling at the bottom of the flask. Starch solution not made properly will appear cloudy.
The Potassium Iodide will ‘go off’ if left for too long in the presence of air, as the iodide ions which it contain will spontaneously oxidise to iodine in the presence of oxygen. This is indicated by a brown-colour of iodine displayed by KI solution. The solution should be made up fresh again when the colour becomes visible to the human eye, in order to prevent deviations in the specified concentration.
Carrying out the reaction. The waterbath was set up at the required temperature. 100 ml of each reagent used (Potassium Iodide, Sodium Thiosulphate, Potassium Peroxydisulphate and Starch Solution) was placed within a beaker in the waterbath, each with its own labelled syringe. 50 ml beakers were left in the waterbath for 5 minutes to allow the beaker itself to be warmed up to the required temperature. After the temperature settles (and this is monitored by several thermometers, one in the waterbath and another in a beaker of distilled water in the waterbath), the following reaction mixture is made up in one of the 50ml beakers at the correct temperature:
1.0 cm3 of Starch Suspension (1%)
2.0 cm3 of Na2S2O3 Solution (0.01M)
3.0 cm3 of KI Solution (0.2M)
The volumes were measured using the syringes in the beakers, as accurately as the syringe will allow. The reagents were added in the order shown. A syringe containing 4.0 cm3 of 0.01M K2S2O8 solution was prepared. This was then injected as quickly as possible into the reaction mixture, and the timer started. There was no need to mix the mixture as the action of squirting in the last reagent actually did the mixing for you. The timer was stopped when the first trace of blue was detectable in the reaction mixture, and the time taken to react noted.
This procedure was repeated many times to give us an idea of the size of the random errors, and by using an average in an attempt to minimise its effect on the final results. Also the procedure was repeated at a range of temperatures between the room temperature (20oC) and 70oC in increasments of 10 oC. A graph of -ln(time) against 1/Absolute Temperature was plotted to determine the activation enthalpy for the particular reaction.
The procedure was repeated when the catalysts were needed. Initially it was decided that the catalyst should be added dropwise after the reaction mixture has been prepared, but before the K2S2O8 was injected. However, the difference in the size of the drops were hampering the consistency of some of the results, hence a new approach was adopted when the practical was being carried out. Instead of using the 0.1M solution, it was decided that a much more dilute solution should be used. The average volume of a pipette drop was calculated using the balance and water (it was 19 ml), and the concentration required if 1 ml was to replace one drop was calculated as 0.0039M. In the end the catalyst was added in the form of a very dilute solution and 1 ml was easier to measure using the existing syringe and much more accurate.
In order to determine the order of the reaction experiments were done at differing concentration. The procedure was repeated but instead of preparing the reaction mixture in the manner specified above, an alternative way of producing a lot of different concentration very quickly was used. The reaction mixture was made using the following recipe:
1.0 cm3 of Starch Suspension (1%)
2.0 cm3 of Na2S2O3 Solution (0.01M)
1.0 cm3 of CuSO4 Solution (0.0039M)
x cm3 of KI Solution (0.2M)
(5.0 - x) cm3 of Distilled Water, and
5.0 cm3 of K2S2O8 Solution (0.01M) to start the reaction
This allowed for a range of concentrations to be produced without making up any more standard solutions. The maximum value of x is kept at 5.0 ml to give a reasonable reaction time.
To vary the concentration of Potassium Persulphate a similar tactic was deployed, with 5.0 ml of KI each time with different amount of distilled water, and injecting complementary amount of Persulphate to start the reaction. As the time it took to inject did not differ by much when compared to the time it takes to complete the reaction, this source of error can be ignored. When the variation in concentration of the catalyst is required, the same method was used again but as 1 ml is a bit more difficult to subdivide with a syringe a teat pipette was used to measure precise amount of the solution in increasments of 0.25 ml.
Practical Considerations
During the actual practical session several difficulties were encountered, and they are addressed here.
Using Starch as an indicator. The mechanism by which starch work as an indicator for iodine is by means of forming a dark-coloured complex with iodine molecules. This reaction in most circumstances can be regarded as instantaneous. When a iodine is produced at a relatively low rate (such as working this particular experiment at 20oC), the solution darkens almost without intermediate stages. As test-runs for this practical were conducted at room temperature, it had not been foreseen at the planning stage that at temperatures above 75oC the starch does not work at all, taking almost 3 seconds to turn fully dark-blue after the yellow-brown colour of the iodine had become visible. Realistically speaking, 70oC is probably the highest temperature at which starch solution could be used as an indicator. It is unclear whether if this time-delay is caused by the high rate at which the iodine is produced, or whether if the solution is affected by the high temperature.
Timing the reactions. I have chosen quite carefully, though a series of test-runs, to find the optimum concentration for the temperature range that I was trying to investigate in the given timescale. There were generally no timing difficulties until the introduction of Copper(II) ions as catalyst. Because of its extremely high efficiency I have had to repeat the experiment with a lower concentration, but the results will still be comparable as we can calculate the activation enthalpy of the reaction by using the gradient of the -ln(time) vs 1/Temperature graph.
Using Copper(II) Ions as catalyst. When 0.1M Cupric Sulphate solution is added to the reaction mixture containing Starch, Sodium Thiosulphate and Potassium Iodide, a light-blue precipitate appears forms and the blue-black colour of the iodine-starch complex appears. After stirring the colour disappears leaving a slightly cloudy solution.
Starting point. As it takes less than one second to squirt the Potassium Peroxydisulphate solution into the reaction mixture to start the reaction, I have decided to define the starting point of the reaction as the start of the squirting operation. Even if one drop gets into the reaction mixture the reaction would have started, so in practise I must squirt as quickly as I can and the one second it takes would be part of the systemic errors present in the set of data obtained.
Reaction Time. The human reaction time has a minimum value of 0.2 seconds (this can be measured using a simple computer programme) but often depends on whether one is anticipating the event. In order to prevent the reaction time from distorting the result during the practicals I must arrange for the timer to be unseen from where I am so that I can’t anticipate when it would go blue. This error can then be quantified as a systematic error of approximately 0.5 seconds.
The order in which the reactants are mixed. By using Cupric ions as catalyst we introduced another problem which I was not aware of until the practicals started: The iodide ions in the solution reacted with the Cupric ions in the solution in significant amounts - enough to alter the results quite dramatically, almost halving the reaction time needed. This is due to the fact that Cupric ions react with any iodide ions present to form a light blue precipitate, and therefore decreasing the concentration of iodide ions. The problem could be solved by changing the order in which the reactants are mixed: if I added the Cupric ions before the addition of Potassium Iodide, but after Sodium Thiosulphate has been added, the Cupric ions will be complexed by the Thiosulphate ions present, and will therefore be prevented from reacting with the iodide as the iodide enters.
Again, as long as within a set of data the order of mixing stays one way, the data still can be compared across data sets because the activation enthalpy is not affected by the concentration of either iodide ions or Cupric ions. It seemed that the relative local concentration of the Cupric ions had an effect, too. When investigating the effect of concentration of the catalyst I had made up some 0.0039M Cupric Sulphate solution, and this did not seem to give the precipitate even though it is injected into the reaction mixture after the addition of iodide ions.
