Chemical Kenetics 17 Exp 3

Chemical Kenetics

A. B. C. Doloiras
Department of Chemical Engineering, College of Engineering
University of the Philippines, Diliman, Quezon City, Philippines
DATE PERFORMED: June 28, 2013
INSTRUCTOR’S NAME: Ms. Irina Diane Castanos

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ABSTRACT This experiment aims to determine the influences of varying factors: temperature, concentration of reactants, and presence of catalysts, on the rate of a reaction. The rate law of a reaction was determined using the initial rates of method. Data gathered from six different runs of a reaction between thiosulfate and hydronium ion was used to plot a linear equation based on the Arrhenius equation. From the equation, the activation energy, Ea, of the reaction was obtained with an experimental value of 58.03 kJ/mol. The experiment showed that higher temperature leads to an increase in the rate of reaction, that a decrease in temperature would do the reverse, and that the concentration of reactants is directly proportional to the rate of reaction. The experiment also showed that a catalyst boosts the rate of reaction as it offers an alternative route of reaction with a lower Ea
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RESULTS AND DISCUSSION

In this experiment, concentration of reactants, temperature, and presence of catalyst, were applied in varied values to determine its effect on the overall rate of the reaction. The influence of the concentration of reactants on the rate of reaction was observed in the first part of the experiment. This was exhibited by the reaction between thiosulfate and hydronium ion in varied molarities. The reaction has the balanced equation

Na2SO3(aq) + 2HCl(aq) → SO2(g) + S(s) + H2O(l) with a net ionic equation, S2O32-(aq) + 2H+(aq) →SO2 (g) + S(s) + H2O(l)­

Using the values of the change of molarity with respect to time, we can express the reaction rate as rate = -Δ[S2O32-]/Δt = -Δ[H+]/Δt
= Δ[SO2]/Δt = Δ[S]/Δt =Δ[H2O]/Δt
As the formation of precipitated sulfur happens, the solution becomes more turbid until it is completely opaque. The opacity indicates the completion of the reaction as it indicates that the limiting reagent is already used up leading to abundant sulfur product precipitation. As a control, an “X” mark was placed on the outer surface of the base of similar, in terms of size, beakers. It was in these beakers where the reaction was left to happen. The time when the mark is no longer visible to the observer indicates the peak or end of reaction; and is recorded as the time of reaction, t.
From the data, the rate of reaction was approximated to be equal to 1/t. This proposition is valid since the reaction stops at a particular and constant molarity, and only the time unit was varied. The only consideration would be on the changes in reaction rates as indicated by the “disappearance” of the mark. Thus, we can take the change in concentration of reactant to be constant leaving us the equation rate = k/t or rate α 1/t.
In total, there were six runs of the reaction done. On runs 1-3, varying volumes of thiosulfate was first mixed with corresponding amount of water, before adding a certain volume of hydrochloric acid which was held constant for these runs. On runs 4-6, varying amounts of HCl and water were mixed together first before adding a constant volume (for these series) of thiosulfate. This made the reagents practically undiluted and can be treated as having a constant concentration in the corresponding run. The concentration of the reagents in consideration is important in determining the rate of reaction (with respect to the reagent).
The reaction order with respect to H+(aq) is zero while the reaction order for S2O32-(aq) is one. The reaction has an overall order of one and has the rate law, rate = k[S2O32-] where k is the reaction constant; and the concentration of S2O32- in the entire solution is represented as [S2O32-].
As shown in the rate law of reaction, the rate of reaction is dependent solely on the concentration of thiosulfate. And since the overall order of reaction is equal to one, it is said that the rate determining step of the reaction to be unimolecular. From these theories, it can be inferred that the reaction follows the mechanism, S2032-(aq) → S(s) + SO32-(aq) slow SO32-(aq) → SO2(g) + ½ O2 fast 2H+(aq) + ½ O2 → H20 (l) fast
Certain intervening variables must be kept constant to keep errors in calculations to the minimum.
First, it is essential that the diameters of the beakers used were kept constant for the opacity of the mixture is dependent to the dimensions of its container. Changes in the dimensions of the beaker would lead to either an increase or decrease of the time it will take for the “x” mark to be covered by the solution as seen by the observer.
The timer used must be the same throughout the experiment because they vary in accuracy. It is important because the reaction rate is measured per unit of time. Changes in accuracy of timer might lead to inconsistencies while solving for the rate of reaction.
In the experiment, instead of using a spectrometer, human perception was used as the basis for the opacity of solution. Opacity in human terms is qualitative and varies from person to person due to eyesight. Thus, it is only necessary that only one observer will judge the opacity for all the runs.
In the second part of the experiment, the influence of the temperature at the beginning of the reaction to rate of reaction was tested. As a control, the concentrations of the reactants were held constant. The reaction was tested in varying temperatures: at 353.15 K, 293.15 K, and 283.15 K. The data shows a trend that the reaction proceeds faster as the temperature where the reaction took place increases. Meaning the rate of reaction is directly proportional to the temperature at which the reaction took place.
This relation between temperature and rate of reaction is caused by the difference in kinetic energy of the reaction. Temperature, as a measure of heat, is equal to the average kinetic energy of an object’s molecules. An increase of temperature would excite the molecules, causing them to move faster and hit each other more frequently, increasing the favorable product-forming collisions, and eventually leading to an increase in rate of reaction. Decreasing the temperature, however, would then lead to a slower reaction. Collision theory can be applied regardless of the thermodynamic states of reactants.
Table 1. Experimental Data from Part B of Chemical Kinetics
Temperature (K)
1/Temp (/K)
Time (s)
1/Time (/s)
353.15
0.00283
6.1
0.164
293.15
0.0034
151
0.00662
283.15
0.00354
1290
0.000775

The results of the experiment are as shown. To be able to create an equation which will allow us to express the relationship between the run time and temperature, we use the Arrhenius equation, k = A e-Ea/RT where k = rate constant; A = Arrhenius constant; Ea = activation energy; R = ideal gas constant; and T = temperature.
Taking the natural logarithm of both sides of the equation, we obtain a linear equation, ln k = - E­a/RT + ln A where - E­a/RT being the slope of the line. Using the argument of k being equal to 1/time of reaction, we arrive at the equation, ln 1/t = - E­a/RT + ln A
Using the data gathered from the experiment, we arrive at the linear equation, y = -6980x +18.08 which has a linearity coefficient of 0.94, which is close to 1. This confirms that the overall order of reaction is 1 since the graph follows a straight line with a negative slope.
The activation energy is the minimum energy required for a reaction to take place. In this case, the activation energy can be readily calculated by multiplying the slope of the line with the ideal gas constant which, in this experiment, is equal to 58.03 kJ/mol. In the experiment, the reaction in consideration is positive, as well as in any other reaction. This trend for the sign of Ea can be justified by the mere meaning of activation energy. It is the energy required to be absorbed by the reactants to break bonds for the formation of products which is true for both exothermic and endothermic reactions. Its sign will always be positive since it is energy absorbed by the system.
The ΔH of the reaction is 65.9 kJ/mol. We observe that ΔH is greater than the calculated activated energy, which is impossible for an endothermic process. However, since ΔH is a theoretical value obtained through thermochemistry, and Ea being an experimental value obtained through chemical kinetics, it is only natural that there may be some errors in calculations or experimental observations. Hypothetically, the reaction has an energy diagram as shown in Fig. 1.

Figure 1. Energy Profile of the Reaction
The last part of the experiment dealt with the influence of presence/absence of a catalyst to the rate of reaction. On the first half of part C, an oxidation reaction as means of reference was done. The reaction is an oxidation of a tartrate ion with hydrogen peroxide, producing carbon dioxide gas and methanoate ion. It has a balanced equation, Na2C4H4O6+3H2O2 → 2CO2 + 2NaHCO2 + 4H2O
On the case in which no catalyst was applied, approximately eight minutes were consumed until the reaction was complete (decolorization). While on the case in which CoCl2, being the catalyst, was added, the pink mixture quickly turns to green after a minute, followed by rapid production of gas bubbles - caused by the oxidation of tartrate ions, yielding considerable amounts of carbon dioxide gas. The green color indicates the formation of Co(III) which is thought to be the actual catalyst, from the oxidized Co(II). Co3+ then bonds to the tartrate ion, allowing oxidation of tartrate ions to take place. At this time, the green solution reverts back to pink, as the Co(III) complex is reduced back to Co(II). The reverting back of color took 30 seconds, having a relative total time of 1.5 minutes for the catalyzed reaction. The cobalt catalyst provides an alternative route for the reaction to occur, which has a lower activation energy; thus, having a faster rate of reaction.
On the second half of the catalysis part of the experiment, a reaction between oxalate and permanganate was done. The reaction has the net ionic equation,
2MnO42-(aq) + 16H+(aq) + 5C2O42- → 10CO2(g) + 8H2O(l) + 2Mn2+(aq)
During the experiment, the time at which the permanganate decolorizes is so small that inconsistent quantitative values or data may be gathered. Under the discretion of the observer, the reactions were just observed repeatedly qualitatively and relative to each other. It is noted that the addition of the first drop of KMnO4 to the oxalate in acidic solution, seemingly took the longest time to decolorize. On the 2nd drop of permanganate, an even faster decolorization was observed. Finally, on the run at which MnSO4 was added before dropping permanganate, the fastest decolorization happened.
This reaction is a example of autocatalysis. An autocatalyst is a product that speeds up its own reaction, being a catalyst itself to the process. In this particular reaction of the experiment, Mn2+ served as an autocatalyst. This is supported by the observation that the permanganate droplet decolorizes faster as the concentration of Mn2+ in the solution increases.

REFERENCES
[1] Chang, R. Chemistry 8th Edition. McGraw-Hill Companies, Inc., 1221 Avenue of the Americas, New York. 2005
[2] Petrucci, R.H., Harwood, W.S., Herring, F.G. General Chemistry: Principles and Modern Applications 8th ed. Pearson Education South Asia Pte. Ltd. Singapore 2004

[3] Whitten, K.W., Davis, R.E., Peck, M.L., Stanley, G.G. Chemistry 8th ed. Thomson Higher Education, California USA. 2007

[4] Brown, T., LeMay, H., Bursten, B., Burdge, J. Chemistry: The Central Science 9th ed. Pearson Education South Asia PTE LTD., 2004.
APPENDIX

See data sheet attached for raw data.

CALCULATIONS:

References: [1] Chang, R. Chemistry 8th Edition. McGraw-Hill Companies, Inc., 1221 Avenue of the Americas, New York. 2005 [2] Petrucci, R.H., Harwood, W.S., Herring, F.G. General Chemistry: Principles and Modern Applications 8th ed. Pearson Education South Asia Pte. Ltd. Singapore 2004 [3] Whitten, K.W., Davis, R.E., Peck, M.L., Stanley, G.G. Chemistry 8th ed. Thomson Higher Education, California USA. 2007 [4] Brown, T., LeMay, H., Bursten, B., Burdge, J. Chemistry: The Central Science 9th ed. Pearson Education South Asia PTE LTD., 2004. APPENDIX See data sheet attached for raw data. CALCULATIONS: