Methodology

METHODOLOGY

The first part of the experiment was done to determine the effect of concentration on the rate of reaction. A beaker, with an X mark at the outer-bottom side, was filled with the required volumes of Sodium Thiosulfate and water. Enough HCl was then added to the solution and the time it took for the X mark to be covered by a cloud of precipitate was recorded. This was done for 3 Runs (runs 1-3; refer to table 1). For the next 3 runs (runs 4-6), required amounts of HCl was added into the water (Refer to table 1). Required volumes of Sodium Thiosulfate were then added to the solutions and the time it took for a cloud of precipitate to cover the X mark was recorded. Take note that only 1 beaker must be used throughout the experiment and that there should be only one (1) observer.Table 1.1 Determination of the Effect of Concentration on the Rate of ReactionRunVol. of 0.15 M Na2S2o3 , mLVol. of water, mLVol. of 3M HCl , mL

110.03.02.0

25.08.02.0

32.510.52.0

45.01.01.5

55.01.51.0

65.02.00.5

Next was the determination of the effect of temperature on the rate of reaction. Specified volumes from Run 3 (Table 1) were used to conduct this part of the experiment. Sodium Thiosulfate and water were mixed together in a 50-mL beaker with X mark while HCl was put in a test tube. Both solutions were put inside a water bath of desired temperature (cold, ambient, hot). It took about 5 to 10 minutes for the temperature to remain constant and this temperature was recorded. After this, HCl was added to the thiosulfate solution. Keep in mind that the 50-mL beaker was still submerged in water bath. The time it took for the cloudy precipitate to cover the X mark was recorded. Waste from the first part and this part was filtered. Precipitate was collected ad submitted in a vial while the filtrate was disposed of in the Inorganic Waste bin.

The last two parts of the experiment dealt with Catalysis. For the first, a hot bath was prepared and its temperature was kept at around 650C. Two 6-inch test tubes were then prepared. Both test tubes contained 5mL of 0.3M sodium tartrate and 2mL of 6% Hydrogen peroxide. An additional 8 drops of 0.3 M CoCl2 was added to test tube 2. After that, both test tubes were placed in the hot water bath and the time it took for the solutions color to change was recorded. Keep in mind that the temperature must be kept around 650C since the reaction mixture may overflow at higher temperatures. Waste from this part was disposed of in the Peroxide Waste container.

For the last part, 1 mL of saturated Na2C2O4, 5 drops of water and 1 mL of 3.0 M H2SO4 were mixed in a 4-inch test tube. This mixture was then split into two, leaving us with Test Tube 1 and Test Tube 2. A drop of 0.01 M KMnO4 was added to test tube 1 and the time it took for permanganate to decolorize was recorded. Another drop of 0.01 M KMnO4 was added to test tube 1 and again, the time it took for permanganate to decolorize was recorded. To test tube 2, 5 drops of 1% MnSo4 was added and then mixed thoroughly. A drop of 0.01 M KMnO4 was then added to the mixture and the time it took for permanganate to decolorize was recorded. Waste from this part was diluted with 0.1M NaOH and was disposed of in the Inorganic Waste bin.

RESULTS AND DISCUSSION

The experiment proceeded without much variation from what should happen according to the procedure. The equation below (Equation 1) presents the balanced net ionic equation of the reaction and the table below (Table 1.2) shows the results from the first part of the experiment.S2O32-(aq) + 2H+(aq) ( SO32-(aq) + S(s) + SO2(g) + H2O(l) (1)Table 1.2 Results for the reaction of thiosulfate

RunTime (s)1/Time (s-1)[S2O32-] initial (M)[H=]initial(M)

125.090.039860.53.0

262.290.016050.93753.0

3108.480.009220.35713.0

431.970.031280.154.5

534.00.029410.152.0

641.290.024220.150.75

It can be observed that a change in molarity of the reactants has happened. This is reasonable considering that there would be an initial reaction of water with Sodium Thiosulfate or Hydrochloric acid.This information is necessary to us because we can identify its rate constant now. The rate can now be expressed as that of the rate law (Equation 2):Rate = k [S2O32-][H+]2 (2) ; where 1/t = rate

We use 1/t as the measure of the initial rate of reaction. This is possible because of the assumption that all our reactions proceed to completion and is composed of the same reactants and produce the same products.[1] What is the substance that causes visibility to change? you may ask. As stated in the methodology, a precipitate causes the clear solution to be cloudy. This is because of the Sulfur solid produced. Although the data seems right, there are still some sources of error in the methods done. The observer might not have been able to start or stop the stopwatch once acid was added. Second is that they mightve had the timer stopped at different concentrations of the cloudiness (although it was stated that there should only be one observer, there will still be some human error from that observer on how he/she perceives the cloudiness).

The second part of the experiment discusses how temperature affects the rate of reaction. A table below is used to present the data gathered:

Table 2 Results for determination of the effect of temperature to rate of reactionTemperature (K)1/Temperature (K-1)Time (s)1/Time (s-1)

303.650.0032977.140.01296

323.150.0030923.430.04268

277.650.00360595.240.00168

Based on the data presented, the rate of reaction increases as temperature increases. According to the collision theory, particles must collide together with enough frequency with sufficient energy. [2] It is known that particles would move faster once they are heated, leaving us with more possibilities of collision from these substances. Relating this to Arrhenius equation, the linearity of ln k versus 1/T is ideally 1. Computing for the regression, we end up having an R2 that is equal to 0.998. The closer to 1 the regression gets, the better the fit of the line is to the points. Relating this to the data, it only means that ln k can be equal to 1/T but not al all times. It is possible that one value might be lower than that of the other. [3] Once this has been computed we can substitute values to the Arrhenius equation to get the amount of activation energy (Ea) needed for the reaction. The Ea of this reaction is positive which is fairly reasonable since rate of reaction increases as temperature increases. In reactions where rate of reaction decreases as temperature increases, the Ea would be negative. So the sign of activation energy not only tells us how much energy is needed to proceed to completion but also whether an increase or decrease in temperature is necessary.

For the first part of the last experiment, there is one method error that I would like to bring up; the temperature of the water bath. As much as we want it to be a constant 650C, sometimes its just not possible to keep it that way for the entire procedure. We would either have to risk decreasing (taking it from the hotplate) or increasing (keeping it in the hotplate) the temperature and the only better solution was to decrease it since an overflow might occur. Moving on to the data, an equation below shows the balanced equation (Equation 3) for the oxidation of tartrate by hydrogen peroxide.

C4H4O62- (aq)+ 5H2O2(aq) + 2H+( 4CO2(g) +8 H2O(l) (3)From the balanced equation, we now know that it is the carbon dioxide that is responsible for effervescence. What gives the color, though, is the addition of CoCl2 forming a cobalt complex first that produces the color pink. The color changes represent how a catalyst creates alternative paths to achieve end point.

For the last experiment, a reaction between permanganate and oxalate occurred. The equation below is the net ionic equation of the redox reaction (Equation 4):

2 MnO4-(aq ) + 5 C2O42-(aq ) + 16H+(aq) --> 2Mn2+(aq ) + 10 CO2(aq ) + 8 H2O (4)Experimentally, it can be proven that the catalyst for the reaction is Mn2+, taking only 6.64 seconds to decolorize the permanganate ion. But for the reaction without the manganese, the second reaction somehow still occurred faster than the first. This is because of autocatalysis. In the initial process, a complex ion containing manganese has been formed so when the next drop of permanganate was added, the reaction proceeded faster. The autocatalyst can be determined by first determining the catalyst and then looking at the possible products of the initial reaction. [4]

Appendix:

1. Plot of ln k vs 1/T

2.

References:

[1] Bissonnette, C.; Herring, G.; Madura, J.; Petrucci, R. General Chemistry: Principles And Modern Applications; 10th ed.; Toronto, 2011.[2] Padolina, M.; Simon-Antero, E.; Alumaga, M. Conceptual And Functional Chemistry; 2nd ed.; Vibal Publishing House: Manila, 2010; pp. 268-269.

[3]People.duke.edu,. Introduction to linear regression analysis http://people.duke.edu/~rnau/regintro.htm (accessed Jun 24, 2015).

[4]Chemeddl.org,. ChemEd DL Image Collection http://www.chemeddl.org/alfresco/service/org/chemeddl/video/video_images?id=vid:691&guest=true (accessed Jun 24, 2015).