Investigating the Effect of Concentration on an Iodide Persulphate Reaction, and Rate aw Determination ab Performed on Monday, February 25 th, 2013
Introduction The purpose of this lab is to observe the effect of concentration on reaction rates, to ultimately determine the rate law for the reaction between the iodide ion (I - ) and the persulphate ion (S 2 O 8 ). The general reaction can be described as (NH 4 ) 2 S 2 O 8(aq) + 2KI (aq) I 2(aq) + 2K 2 SO 4(aq) However, we know that in aqueous solution, the ions will dissociate (except for iodine) into their ionic components. The ions that are not a part of the actual reaction do not affect the ionic compounds in any ways, as they are unreactive. The reaction can be rewritten as: S 2 O 8 (aq) + 2I - (aq) I 2(aq) + 2SO 4 (aq) In this experiment, the experimenters are interested in using the rate of oxidation of the iodide ion by the persulphate ion to determine the rate law for this reaction. The rate of reaction between the two ions will be measured near the beginning of the reaction, before the concentrations of the components in the system have changed significantly. New trials will be run such that one concentration of a specie is altered from that of the original reaction, and the rate will be measured once again. To determine the reaction rate, the amount of product formed or reactant used within a time interval must be measured. For this experiment, the time required for a specific amount of diiodine (I 2 ) to be produced will be measured. To monitor when a specific amount of diiodine has been produced, the diiodine will have to react with a thiosulphate ion. I 2 (aq) + 2S 2 O 3 (aq) 2I - (aq) + S 4 O 6 (aq) As long as thiosulphate ions are still present in the aqueous solution as the reaction proceeds, the reaction will remain colourless (in the presence of starch) since the thiosulphate ions will be using up the diiodine molecules. According to the chemical reaction formula, two moles of thiosulphate is required to react with one mole of diiodine. The first appearance of free diiodine is detected by its reaction with starch. This means that the initial amount of thiosulphate ions added into the reaction has been used up, so there is nothing for the diiodine molecules to react with other than starch. The reaction between diiodine and starch will yield a violet/dark-blue colour due to the starch-iodine complex formed. The reaction should be timed from the moment the two solutions are mixed together to the moment a violet/dark-blue colour is visible. Therefore, the amount of iodine produced within this time interval will be known since the amount of thiosulphate used is known. The process of thiosulphate ions reducing iodine back to iodide before the iodine can form its complex with starch to yield a violet colour is called a clock reaction because this process is used to clock the reaction between persulphate and iodide. 2 S 2 O 3 Average Rate of the Reaction= 1 2 To determine the order of reaction with respect to the iodide ion, the concentration of persulphate must remain constant as the concentration of iodide ion varies (evine, 2002). The relationship between the change of concentration of iodide ion to the time it takes for a constant amount of thiosulphate to be used up will aid the experimenters in determining the order of the reaction with respect to I -. (evine, 2002). 2
In the primary iodide persulphate reaction, we are dealing with two negative ions. This means that the ionic strength of the reaction must be considered when determining the rate of reaction or rate law because the rate of reaction increases as the ionic strength of the reaction increases. The reaction rate constant therefore is dependent on the ionic strength of the reaction (on top of temperature, pressure and other various conditions), so it is necessary to maintain a constant ionic strength throughout this can be done by adding an electrolyte such as (NH 4 ) 2 SO 4 or KNO 3. The equation for calculating ionic strength is as follows: μ= 1 where C i is the molar concentration 2 2 C i Z i of the i th species, and Z i is its ionic charge. The ionic strength is a measure of the overall concentration of the ions (evine, 2002). The ionic strength of a solution decreases as the concentration of ions decrease and their valence or charge decreases (Gillenwater, 2002). Experimenters will use their experimental data to observe the relationship between reactant/product concentrations and the reaction rate to deduce the rate law of the reaction between the iodide ion and persulphate ion. Materials The materials used for this experiment is outlined in the CHEM 123 lab manual, Experiment #1. All materials were used without deviation. Experimental Procedure The experimental procedure used for this experiment is outlined in the CHEM 123 lab manual, Experiment #1. The actual experiment deviates from the experimental procedure such that only 6 runs were performed instead 7 runs. Experimental Observations Table 1: Original Observations Preparation of Stock Solutions (NH4)2S2O82 (0.100M) KI (0.200M) Na2S2O3 (0.0100M) Mass of Solid Added (+/- 0.01g) Volume of dh2o added (m) 25 25 25 250 250 250 Run # Table 2: Qualitative Observations Rate Measurements using Varying Concentrations of Reactant and Product Add to Erlenmeyer 0.1M (NH 4 ) 2 S 2 O 8 m 0.1M (NH 4 ) 2 SO 4 m 0.2M KI m 0.2M KNO 3 m Add to Beaker 0.01M Na 2 S 2 O 3 m H2O m 0.2% Starch m Elapsed Time 1 20 0 20 0 10 0 5 5:03 2 10 10 20 0 10 0 5 9:27 3 5 15 20 0 10 0 5 22:08 3
4 20 0 10 10 10 0 5 9:25 5 20 0 5 15 10 0 5 16:46 6 10 0 20 0 10 10 5 10:32 7 20 0 5 0 10 15 5 23:00 Results and Calculations Below shows a table indicating all of the results and calculations for runs #1-7 Run # [S2O8] [I-] [S2O3] S2O8 t Rate ([ S2O8]/s) Rate Constant (k) Ionic Strength 1 0.0364 0.0727 0.00182 0.000909 5:03 0.000003 0.00113 0.187 2 0.0182 0.0727 0.00182 0.000909 9:27 0.0000016 0.00121 0.133 3 0.00909 0.072 0.00182 0.000909 22:08 0.000000685 0.00104 0.105 4 0.0364 0.0364 0.00182 0.000909 9:25 0.00000161 0.00122 0.151 5 0.0364 0.0182 0.00182 0.000909 16:46 0.000000904 0.00136 0.133 6 0.0182 0.0727 0.00182 0.000909 10:32 0.000001438 0.00109 0.133 7 0.0364 0.0182 0.00182 0.000909 23:00 0.000000659 0.000995 0.133 Note: calculations will be sample calculations from trial #1 for the concentrations of S 2 O 8, I -,S 2 O 3 and - S 2 O 8 Note: Total volume in reaction vessel for all trials is 55m Calculating the Concentration of S 2 O 8 Using 20m of 0.1M (NH 4 ) 2 S 2 O 8(aq) in trial #1 V 1 C 1 =V 2 C 2 (20m) (0.1M)=(55m)C 2 C 2 =0.0364M Therefore the concentration of S 2 O 8 (aq) in trial #1 was 0.0364M. Calculating the Concentration of I - Using 20m of 0.2M KI (aq) in trial #1 V 1 C 1 =V 2 C 2 (20m) (0.2M)=(55m)C 2 4
C 2 =0.0727M Therefore the concentration of KI (aq) in trial #1 was 0.0727M. Calculating the Concentration of S 2 O 3 Using 10m of 0.01M Na 2 S 2 O 3(aq) in trial #1 V 1 C 1 =V 2 C 2 (10m) (0.01M)=(55m)C 2 C 2 =0.00182M Calculating the Change in Concentration of S 2 O 8 from the start of the reaction to when a violent/darkblue complex is formed - S 2 O 8 can be calculated by dividing the initial concentration of thiosulphate (S 2 O 3 ) by two because two moles of thiosulphate are used up per mole of iodine. This is evident in the chemical equation below. I 2 (aq) + 2S 2 O 3 (aq) 2I - (aq) + S 4 O 6 (aq) The amount of iodine produced will be half the amount of thiosulphate ions added. The amount of iodine produced is also indicative of the amount of S 2 O 8 left in the reaction vessel (1:1), which will enable us to calculate the change in concentration of S 2 O 8. Moles of thiosulphate Added Moles of Iodine Produce d= 2 0.00182 2 0.000909 moles Moles of Iodine Produced = Moles of Persulphate Ions Produced S 2 O 8 2 = (0.000909 ) S 2 O 8 2 =0.000909moles Calculating the Time it took for the Reaction to use up all the Thiosulfate Ions For trial #1, t = 9:27 minutes = 567s Calculating the Rate of the Reaction (for Trial #2) To calculate the rate of the reaction, the rate of consumption of persulphate ions must be calculated. To do this, the change in concentration of persulphate ions (S 2 O 8 ) will be divided by the time it took for the reaction to use up all the thiosulphate ions ( t), which is essentially the time it takes for the reaction to produce a certain amount of iodine, which in turn gives the time it takes for a certain amount of persulphate ions to be converted into products ( S 2 O 8 ). 5
Rateof Reaction= S 2O 8 2 0.000909M 567s 0.0000016M/ s t Calculating the Rate Constant (k) I Rate= S O 2 8 t 2 =k[s 2 O 8 2 ] m k = rate constant m = order of the reaction with respect to S 2 O 8 n = order of the reaction with respect to I - To find the value of m, a graph of log( t) vs. log[s 2 O 8 ] will be plotted, using values from runs 1, 2, 3. From the graph, the slope will be calculated, and from there, the order of the reaction with respect to S 2 O 8 can be determined. The equation for the linear trendline of this data is 1.0651. y=1.0651x 0.9327 Therefore, the order of the reaction with respect to S 2 O 8 is 1, first order. m = 1, indicating that the slope is To find the value of m, a graph of log( t) vs. log[i - ] will be plotted, using values from runs 1, 4, 5. From the graph, the slope will be calculated, and from there, the order of the reaction with respect to I - can be determined. The equation for the linear trendline of this data is 0.8665. y=0.8665x 1.4984 Therefore, the order of the reaction with respect to I - is 1, first order. I Rate= S O 2 8 t 2 =k[s 2 O 8 2 ] m n = 1, indicating that the slope is Rate for trial #2 = 0.0000016 mol - s -, m = 1, n = 1 6
I Rate=k[ S 2 O 8 2 ] m 0.0000016 mol 1 mol =k [0.0182 s ] [0.0727 mol 1 ] 0.0000016 mol mol2 =k [0.00132314 ] s 2 k =0.001209244675445 k=0.00121 mol s Calculating the Ionic Strength for the Solution (for Trial #2) ionic strength=μ= 1 2 C i Z i 2 μ= 1 2 ( [ NH 4 ] (+1 ) 2 )+([ S 2 O 8 ] ( 2) 2 )+ ([ K ] (+1 ) 2 )+ ([ I ] ( 1) 2 )+([ Na ] (+1) 2 )+( [S 2 O 3 ] ( 2 ) 2 ) μ= 1 2 ([ 0.0182 2 ] (+1 )2) + ([0.0182 ] ( 2 )2 )+( [0.0727 ] (+1 ) 2 )+ ([0.0727 ] ( 1 ) 2 )+( [0.00182 2 ] (+1 ) 2 )+( [0.0018 Questions 1. (a.) Calculating the average rate constant for this reaction using data from trials 1 to 5. Average Rate Constant= 0.00113+0.00121+0.00104+0.0122+0.00136 =0.001192=0.00119 5 (b.) The rate constant values from trials 6 and 7 are not used when calculating the average rate constant, because in the solutions made for trials 6 and 7, the species added into the beaker (which is essentially the amount of iodide ions added into the solution) is diluted with 10m and 15m of water respectively. This means that there is less iodide to be reacted with the persulphate ions. 2. (a.) k-value for run #2 = 0.00121 k-value for run #6 = 0.00109 The k-value for run #2 is larger than that for run #6 by 0.00121. There is a difference between these two values because in run #2, 10m of (NH 4 ) 2 SO 4(aq) was added to the Erlenmeyer flask, whereas in run #6, no volume of (NH 4 ) 2 SO 4(aq) was added. The significance of adding (NH 4 ) 2 SO 4(aq) into the reaction mixture is that we are adding the conjugate base of (NH 4 ) 2S2O8(aq) in it. This means that in run #2, since the 7
conjugate base is already added, less amount of (NH 4 ) 2 S 2 O 8(aq) will have to convert into (NH 4 ) 2 SO 4(aq) on the contrary in run #6, since none of the conjugate base was added initially, more (NH 4 ) 2 S 2 O 8(aq) will have to convert into (NH 4 ) 2 SO 4(aq). This will result in larger rate constant for run #2 since there will be some concentration of conjugate base to divide by the persulphate acid. (b.) k-value for run #5 = 0.00136 k-value for run #7 = 0.000995 The k-value for run #5 is larger than that for run #7 by 0.000365. There is a difference between these two values because in run #7, 15m of water is added to the reaction mixture, meaning that everything is diluted, whereas in run #5, no volume of water was added. Diluting the reaction with water means that the reaction will occur less rapidly since the concentration of all the species will be lowered in the presence of water volume as a result, the rate of reaction will be slower, and the rate constant will be smaller as well. This is evident as the rate constant in run #5 is 0.00136 and in run #7 is 0.000995. The rate constant for run #5 is larger as a of less diluted reactants. One factor that affects the rate of reaction, and in turn the rate constant, is the concentration of reactants or products. The higher the concentration, the more particles are able to collide, and thus will result in a quicker rate of reaction. 3. (a.) If the concentration of thiosulphate was doubled, keeping all other factors the same, there would be no effect on the actual reaction itself because thiosulphate in this case is just acting as an indicator for starch. The more concentration of thiosulphate added in the iodide-persulphate reaction, the more time it will take for a violet/dark-blue colour to appear in the solution since there are more thiosulphate ions that can react with the iodine ions. The rate of reaction would be unchanged. (b.) If the concentration of persulphate ions was halved, keeping all other factors the same, the rate of reaction would decrease by half since there are less persulphate ions per liter to react with the iodine ions and since the concentration of persulphate is first order with respect to the whole reaction. Specifically, the rate of reaction would be half of its original value. However, the rate constant will remain unchanged because the same reactants are being used. (c.) If the temperature of the solution was increased, both the rate of reaction and rate constant would increase. The rate of reaction will increase as temperature increases because having a higher temperature correlates to an increase in the average kinetic energy of the reacting system. Having a higher kinetic energy means that a greater portion of the particles will have enough energy to overcome the barrier (activation energy), and thus more reactants will be readily able to convert into products. The rate constant would increase with a temperature increase as well because the rate constant is dependent on the temperature of the system. This can be seen in Arrhenius s equation a formula for the temperature dependence of reaction rates. E a RT k=ae 8
where k is the rate constant, A is the exponential factor, E a is the activation energy, R is the gas constant and T is the temperature in Kelvin. This means that the rate constant changes with respect to any temperature change. 4. (a.) If the concentration of thiosulphate was doubled, keeping all other factors the same, there would be no effect on the rate constant of the reaction because the rate constant is only dependent on the temperature of the environment, ionic strength of the solution and activation energy of the reaction. (b.) If the concentration of iodine ions was halved, keeping all other factors the same, there would be no effect on the rate constant because the rate constant is only dependent on the temperature, activation energy and ionic strength of the reacting solution. The rate constant is only concerned with the orders of different reactants with respect to the reaction. A decrease in persulphate ions will only decrease the number of collisions with iodine per unit time, thus changing the rate of reaction. The rate constant is unchanged. (c.) If the ionic strength of the solution was decreased, the rate constant would decrease as well. The ionic strength of a solution is measure of the strength of the negative ions involved a solution influencing the rate of reaction. As the concentration of negative ions in a solution increases, the ionic (negative charge) decreases as well because the positive ions (spectator ions) that travel around in sync with the negative ions are present, which shields part of the negative charge. If the ionic strength of the solution was decreased, that would mean there is a greater concentration of one of the two ions or both. Decreasing the ionic strength of a solution would decrease the rate constant by increasing the effective activation energy. Discussion Why we used different runs to find m and n Reacting different concentrations of persulphate ions and iodine ions, and observing their different rates of reactions, the rate law for the iodide-persulphate reaction could be deduced as being: 2 S 2 O 8 I Rate aw=k To determine the reaction rate, the amount of product formed or reactant used within a time interval must be measured. For this experiment, the time required for a specific amount of diiodine (I 2 ) to be produced will be measured. Thiosulphate ions, which will be added to this solution will form a little reaction between itself and the iodine product molecules of the original reaction that is being tested. Thiosulphate will act like an indicator for starch in such a way that 2 moles of thiosulphate ions will consume an iodine molecule, and once all the thiosulphate ions have been used, then the starch will start binding to the iodine molecules to yield a violet/dark-blue colour in solution. If we time the reaction from when all the species are mixed to together until a change in colour is visible, we will be able to find the rate of the reaction with respect to how many iodine molecules have been produced. Once the amount of iodine molecules produced is known, the amount of persulphate ions that have been consumed into products will be known from the chemical equation of : S 2 O 8 (aq) + 2I - (aq) I 2(aq) + 2SO 4 (aq) It can be seen from the stoichiometric coefficients that for every iodine molecule (I 2 ) produced, one persulphate ion (S 2 O 8 ) is consumed (1:1 ratio). This means that the rate of reaction can be easily determined for this reaction if thiosulphate is used. 9
To find m, the order of the persulphate ion concentration with respect to the rate of the reaction, only data from runs 1, 2 and 3 were used. The reason for this is because to find the order with respect to one reactant, that specific reactant concentration must vary while all the other species are held constant. In runs 1, 2 and 3, the concentration of iodide ions was held constant, while the concentration for persulphate ions were varying (with its conjugate base of sulfuric acid). When -log t was plotted against log[s 2 O 8 ], the data yielded a linear regression of slope 1.0651. Rounding the slope to the nearest whole number as 1 was indicative of the order of the reaction with respect to the persulphate ion. Persulphate ions being first order means that the rate of reaction is directly proportional to the concentration of persulphate ions if the concentration was doubled, the rate of reaction would double as well. To find n, the order of the iodide ion concentration with respect ot he rate of the reaction, only data from runs 1, 4, 5 were used because in these trials, the concentration of persulphate ions was held constant while the concentration of the iodide ions varied. When -log t was plotted against log[i - ], the data yielded a linear regression of slope 0.8665. Rounding the slope to the nearest whole number as 1 was indicative of the order of the reaction with respect to the iodide ion. Iodide ions being first order also means that the rate of reaction is directly proportional to the concentration of iodide ions. The rate law for the iodide-persulphate reaction was found from one source: 2 S 2 O 8 I Rate aw=0.0040 (The Rate of an Iodine Clock Reaction) From this rate law above, it shows that the value of k is 0.0040, which is different than the experimentally k-value found from this experiment being (average of k-values from runs #1-7) 0.o0115. experimental value theoretical value Percentage Error= theoretical value 100 Percentage Error= 0.00115 0.0040 0.0040 100 Percentage error=71.25 Comparing our experimentally determined k-value to the one found from another source reveals that these 2 values have a 71.25% error. This is a large discrepancy, meaning that there were definitely some errors that were made during this lab. Another factor like temperature could have been the cause of the discrepancy. However, the orders of each reactant seem to agree with each other. In this experiment, the idea of using one reaction to find the rate of another reaction was reinforced a clock reaction. This was possible using starch as the indicator for the presence of iodine ions. The thiosulphate ions would consume the iodide ions, leaving the solution colourless. As soon as the thiosulphate ions are all used up, the solution will have free iodine molecules, forming an iodine-starch complex thus yielding the violet colour in solution. Conclusion Ultimately, the experiment was successful in calculating the rate law for the iodide-persulphate reaction. The experimenters observed the effect of concentration reaction rates such that as the 10
concentration of a reactant increases, the rate of reaction increases, and as the concentration of a reactant decreases, the rate of reaction decreases. The experiment helped carry out the different runs of the iodidepersulphate reaction using different concentrations of each ion. The reaction was timed when 0.00182 moles of thiosulphate was used up, and this was possible in using a clocking technique. ogarithmic graphs were plotted to calculate the orders of the reaction with respect to each reactant specie. The rate law was then determined to be: Rate = 0.00115[I - ][S 2 O 8 ]. If this experiment was redone, to calculate a more accurate and precise rate constant value, each run should be tested for multiple times so that an average t value could be calculated. This would yield a more accurate rate constant because then the experimenters would have an average reaction time to work with, thus taking the value that lies in the middle of the dataset. With only one reaction time value, it is hard to tell if the information agree with each other or are of expected values. One type of random error that could have affected the data is a varying temperature change throughout the period of the lab. If the temperature was not constant all through the lab, the data would have been skewed depending on whether the temperature was decreasing or increasing. Another error could have arose from the fact that the persulphate ions have a small tendency to react with the thiosulphate ions, thus providing an incorrect reaction rate and rate constant. The iodide ion could have reacted with the iodine molecule to produce a triiodide ion, which could have reacted with the persulphate ion. This is also another possibility from which discrepancies could have arisen. References Gillenwater, J. Y. (2002). Adult and Pediatric Urology. Philadelphia: ippincott Williams & Wilkins. evine, I. N. (2002). Physical Chemistry, 5th Ed. New York: McGraw Hill. Moya, M.., Izquierdo, C., & Casado, J. (1991). Microemulsions as a medium in chemical kinetics; the persulphate-iodide reaction. The Journal of Physical Chemistry, 6001-6004. The Rate of an Iodine Clock Reaction. (n.d.). Retrieved March 8, 2013, from comcast: http://home.comcast.net/~drsamples/chemweb401/iodineclockrxnf10.pdf 11