INETICS CHEMICAL K INETICS
P.D. B ATA ELCHOR H ALL , C OLLEGE OLLEGE OF E NGINEERING NGINEERING M ELCHOR , DILIMAN QUEZON CITY , PHILIPPINES UNIVERSITY OF THE PHILIPPINES DATE PERFORMED: DECEMBER 3, 2013 INSTRUCTOR’S NAME : IRINA DIANE CASTAÑOS
ABSTRACT The experiment for Chemical Kinetics aimed to show students how the concentrations of the reactants and the temperature of the environment can affect the reaction rates. Also, it taught the students how to compute for rate laws of experiments, order of reactions, and the concept of catalysts. The experiment demonstrated to the students the abstract concepts to give them a clearer understanding of the topic. The experiment had three parts. Part 1 was carried out by mixing sodium thiosulfate (Na 2S 2O3 ) with hydrochloric acid (HCl) in different concentrations. The time it took for the at the bottom of the beaker, to be hidden by the milky white product was taken. Part 2 was carried “x” at out in a similar way as part 1, the only change being the reactions had to be done in different temperatures. Part 3 explored the concept of catalysts. Two test tubes with similar content (save for the catalyst in test tube 2) were reacted with the same reactant and students observed how much faster the reactions took with the catalysts present. In the experiment, it was concluded that: the reaction proceeds faster the more concentrated the reactants are, the reaction proceeds faster with higher temperatures, and certain substances and solutions can act as catalysts for certain c ertain reactions.
RESULTS AND DISCUSSION
The experiment explored the factors that affect the rates of reactions: the concentration of reactants, the temperature of the environment, and present catalysts. In part one, the concentrations of thiosulfate and hydronium ions were altered to test the effect of concentration on reaction rate. The balanced reaction is as follows: Na2S2O3(aq)+HCL(aq)2NaCl(aq)+H2O(l)+SO2(g)+ S2(s) With the balanced net ionic equation being: S2O32-(aq) + 2H+(aq) → SO2(g) + S(s) + H2O(l). Using the balanced net ionic equation, we can get the reaction rate using the reactants and products. The formulas are as follows: Rate
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The reaction’s was considered complete when the “x” mark at the bottom of the
beaker was covered due to the cloudy white appearance of the produced liquid afterwards, which was caused by the production of S (s). To minimize the amount of errors in this part of the reaction, person that was tasked to time the reaction time and the beaker size were kept constant. In addition, the beaker size used, which was a 50-mL beaker, was chosen for the experiment because of its small diameter, which yielded more collisions of the reactants. The time it took for the reaction to be “complete”
was
recorded
and
its
reciprocal was used as the initial rate of the reaction.
This was a valid method because the absolute mass of sulfur formed was unknown and the initial rate is taken as a measure of the relative rate of the reaction. Also, the initial concentrations of the reactants were constant.
Run No.
Time (s)
1/Time (s-1)
2[ S2O3 ]
[ H ]
1
39.4
0.0254
0.12 M
3M
2
70.5
0.0142
0.06 M
3M
3
171.2
0.00584
0.03 M
3M
4
33.5
0.0299
0.15 M
1.8 M
5
43.4
0.0230
0.15 M
1.2 M
6
52.6
0.0190
0.15 M
0.6 M
+
The data shows that the reaction rate increases when the concentration is increased. This means reactants’
concentration
is
2-
constant and [S2O3 ] is the concentration of the thiosulfate ion. From the experimental rate law obtained, we can have a proposed mechanism composed of elementary steps. Given that the reaction is a first order reaction, the molecularity of the reaction is unimolecular. With that, a possible mechanism for the reaction can be:
Table 1. Part 1 Experiment Results
that
2-
Rate = k [S2O3 ], where k is the rate
directly
proportional to the rate of the reaction. To be able to obtain the rate law of the reaction, the order of the reaction with respect to the thiosulfate and hydronium ion had to be calculated. Using the formula: 2-
Rate = k [ S2O3 ]m [H+]n And with some manipulations of the equation, the following formula can be obtained: [ ] [ ] [ ] [ ]
By using the same concentrations of one ion, we can get the order of the reactions with respect to the other ion can be obtained. The data showed that the orders of the reaction with respect to the thiosulfate and hydronium ion were 1 and 0 respectively, giving the following rate law:
2H+ + ½O2 H2O (fast step) 2S2O3 S(s) + SO32(slow step) 22SO3 ½O2 + SO2 (fast step) 2+ S2O3 (aq) + 2H (aq) → SO2(g) + S(s) + H2O(l). In part two, the same reaction was used and the same instructions were given. The only difference with the part two is that the change was in the temperature of the environment of the reactions: in a hot bath, room temperature, and an ice bath. The following data were obtained: Table 2. Part 2 Experiment Results Temperature (K)
1/Temperature (K -1)
Time (s)
1/Time (s -1)
276.15
1/276.15
948
1/948
299.15
1/299.15
160
1/160
341.15
1/341.15
25
1/25
The data shows that the reaction rate increases when the temperature is increased. This means that temperature is directly proportional to the rate of the reaction. A possible explanation to this relationship could be because the higher the temperature of the environment is, the higher the energy the reactants can obtain, therefore increasing the reaction rate.
The Ea = 43.33 kJ/mol. In this experiment, the reaction in has a positive sign. This trend in the sign of the Ea is because of the requirement for reactants to absorb energy in order to break bonds for the formation of products which is true for both exothermic and endothermic reactions. Hence, all Ea values are positive. Also, with the ∆Hrxn=65.9 kJ/mol, the energy profile of the reaction is as follows:
Figure 1. Graph of ln 1/t vs 1/T
-3.2 -3.7 e-4.2 m-4.7 i t /-5.2 1 n-5.7 l -6.2 -6.7 -7.2
0
2
4
1/Temperature
Figure 2. Energy Profile of Reaction
The figure above shows the graph of the data obtained through the experiment. The equation of the line is y= -5211.64x + 12.44, and the linearity coefficient is 0.9902. The Arrhenius equation can be used to determine the Ea. The equation being:
Where: k = rate constant A = Arrhenius constant Ea = activation energy R = ideal gas constant T = temperature.
Manipulating the formula a bit, we can put it in a form so that it will look similar to the slope intercept equation (y=mx+b) of the graph:
Where: ln k = y
ln A
=m =x =b
Using this, we can determine the activation energy, by multiplying -5211.64 by 8.314 then dividing the answer by 1000 to convert to kJ.
Part three of the experiment explored the concept of catalysts in reactions to speed up the reaction rate. Part three is divided into two subparts: the oxidation of tartrate by hydrogen peroxide (deals with a catalyst) and the reaction of oxalate with permanganate (deals with the an autocatalyst). The balanced equation for the first subpart, which was the oxidation of tartrate by hydrogen peroxide, is as follows: C4H4O6+5 H2O2 4 CO2+6 H2O+OH-(aq) The color change in test tube 1 was from colorless to yellow in 8 minutes. The color change in test tube 2, the one with CoCl2, changed from colorless/red to green in 5:16 minutes. Thus concluding that the catalyst in the first part of part three is CoCl2 since catalysts speed up reactions. Also, the CO2 produced caused the effervescence observed.
In the second subpart, the concept of autocatalysts was explored using the reaction of oxalate with permanganate. The reaction is as follows: 16H++2MnO4-+5C2O422Mn2++10CO2+8H2O
An autocatalyst is defined as an event wherein the catalyst of the reaction comes from one of the products of the reaction itself but with an appropriate reaction substrate. In this reaction, it may have been observed that the reagent species may have already been in the reaction beforehand but still reacted slowly. This shows that the reaction substrate for autocatalysis is not yet set. In this case, seeding of the catalyst needed to be done which was in the form of addition of more permanganate. Through that, the reaction was pushed forward as the reactants produced the catalyst Mn2+ which in turn increased the rate of reaction continuously by forming additional catalyst throughout the process. REFERENCES
[1] Petrucci, Ralph, Geoffrey Herring, Jeffry Madura, Carey Bissonnette (2011). General Chemistry – Principles and Modern Applications. 10th edition. Pearson Education, Canada, 2011. [2] McMullin, Barry (1999). Some Remarks on Autocatalysis and Autopoeisis. Dublin City university, School of Electronic Engineering Artificial Life Laboratory: University of Ghent, Belgium. Technical Report No. bmcm9901. [3] Silberberg, Martin (2007). Principles of General Chemistry. The McGraw-Hill Companies, Inc., 1221 Avenue of the Americas, New York, NY 10020.
