Analyzing Concentration vs. Time Data
Introduction
The area in chemistry designated for studying chemical reaction rates is called chemical kinetics. The reaction rate refers to the change in concentration of a reactant or product in a given time interval, or the rate of reaction is the speed with which either a reactant is transformed or a product is formed, and units are usually molarity per second (M/s).
The reaction rate depends on various factors such as the nature of the reactant, concentration, volume, pressure, temperature and the use of catalysts. In most of the reactions the rate depends directly on the concentration of the reactant, and therefore, the higher the concentration, the faster the rate of the reaction will be and the faster the reaction will occur. This is explained by the collision model, because having a greater number of reactant molecules and concentration will increase the number of molecular collisions which are responsible for the occurrence of chemical reactions.
Objectives
1. Observe the effect of concentration on reaction rates.
2. Develop a mathematical relationship between concentration and reaction rate.
3. Determine the order of a reaction.
4. Observe the effect of a catalyst.
Lab 12: Analyzing Concentration vs. Time Data
Background
A. The effect of concentration
The rate of a reaction is the change of amount or concentration per unit of time and can be measured by observing a reactant or a product. The reaction rate is affected by the concentration, the temperature, particle size, the nature of the reactants, and the presence or absence of catalysts. In this experiment you will look at the effects of concentration.
The collision rate (the number of collisions per second) increases with concentration. The greater the concentration, the greater the collision rate of the reacting ions and thus the greater the reaction rate.
Let’s use an example of people in a shopping mall. Imagine 200 people walking randomly in a shopping mall. Since ions are blind and deaf, the people would need to be blindfolded and required to not make any noise except for an occasional thud as they hit each other once in a while. Now suppose there are 2000 people in the same shopping mall under the same conditions. There would be many more thuds of people bumping into one another. The situation with ions is analogous.
You will use HCl as a catalyst in this experiment. A catalyst is a substance that increases the reaction rate without being consumed. An example of a catalyst is the road salt that is sometimes used when ice forms on a highway. This salt does not cause corrosion on our road vehicles, but does greatly increase the corrosion rate.
In this experiment we will compare the reaction rates of several solutions of sodium thiosulfate and hydrochloric acid. The equation for this reaction is:
Na2S2O3(aq) + HCl(aq) ? S(s) + Na2SO3(aq) + HCl(aq) (Equation 12.1)
The reaction rate will change as the concentration of sodium thiosulfate changes. We will observe the sulfur that precipitates as a yellowish-white powder to estimate the reaction rate. Note that HCl is a catalyst and is not consumed.
B. Reaction order.
The rate of a chemical reaction usually varies with the concentration. This can be expressed by the Rate Law or the Rate Equation:
Rate = k [A]a[B]b[C]c … (Equation 12.2)
where [A], [B], [C], etc are the concentrations of reactants A, B, C … and a, b, c … are exponents. The values of the exponents (a, b, c …) are determined experimentally. The order of a single reaction is the exponent (a, b, c …). The overall order of the reaction is the sum of the exponents. From algebra, when similar terms are multiplied, the exponents are added. In this case the addition of the exponents is not algebraically legal because [A], [B], [C]… are seldom equal to each other. We will ignore this “little” detail.
For an order of zero, the concentration of a reactant has no effect on the reaction rate. Let’s use math to explain this. The form will be:
R = kM0 (Equation 12.3)
where R is the reaction rate, k is a constant and M is the concentration.
Since M0 = 1, the value of R will not vary with M. For zero order the curve of R vs M is a horizontal straight line, as seen in Figure 12.1.
If the order is one, the curve of the rate against the concentration is linear. The form for this is:
R = kM1 (Equation 12.4)
We know that usually an exponent of one is not written. The curve for a first order reaction is linear with a positive slope, as in Figure 12.2.
The curve for order two is parabolic, as seen in Figure 12.3. The algebraic form is:
R = kM2 (Equation 12.5)
Materials
• Reagents:?• 0.1 M hydrochloric acid, HCl?• 0.1 M sodium thiosulfate, Na2S2O3
• 96-well reaction plate • Plastic toothpicks?• Small beaker?• Distilled water
• Stopwatch or a watch with a second hand • White paper
Procedure
Note: the rate of the reaction in this experiment is very temperature-depen- dent. If the solutions are cold, the reaction will be slow. You can speed up the reaction by warming the solutions and the distilled water either in your hands or in warm (not hot) water before you start.
1. Put four X’s on a sheet of white paper so that they will be directly under four adjacent wells in the reaction plate. See Figure 12.4. Put the reaction plate on the paper so that the X’s are visible through the bottoms of the wells in row A. Draw a line around the reaction plate to help you restore the positions of the plate and paper in case you bump or move them.
2. Put five drops of Na2S2O3 in well A1, four drops in well A2, three drops in A3, and two drops in well A4. Record this in the table in your lab workbook.
3. Put one drop of distilled water in well A2, two drops in A3, and three drops in A4. Each well should now contain 5 drops of liquid.
4. Stir the wells with a plastic toothpick. Wash the toothpick in distilled water and shake off the excess water between each well.
5. Start your stopwatch or timer and at 15 second intervals, put three drops of Figure 12.4 hydrochloric acid, HCl, in each of the wells. So well 1 will be -0 seconds, well 2 is -15 seconds, well 3 is -30 seconds, and well 4 is -45 seconds from the final time.
6. Watch the wells and record the times in Table 12.1 when each of the X’s are no longer visible through the precipitate. Remember to subtract the start time difference.
7. Clean the reaction plate immediately. The sulfur precipitate will be difficult to remove if you delay.
Develop Your Own Procedure
Review the data and graph you generated for your investigation on the relation of concentration. In the first part of the experiment, you diluted the sodium thiosulfate to change the concentration but kept the concentration of the hydrochloric acid constant. What do you suppose would happen if you keep the concentration of the sodium thiosulfate constant and varied the concentration of the hydrochloric acid?
Design and perform a procedure to show haw change in the concentration of hydrochloric acid will affect the reaction its rate.
1. Write your procedure, paying attention to details.
2. Set up your data chart.
3. Process your results by performing the necessary calculations and then graphing. Explain how you will do this.
4. Determine the order of reaction.
Post Lab Questions
Table 12.1
Well Drops Na2S2O3 Molarity (M) Time to Cloud (t) (s) 1/t (s-1)
A1 5 0.025M 23s 0.09s-1
A2 4 0.02M 58s 0.04s-1
A3 3 0.015M 150s 6.7×10-3s-1
A4 2 0.01M 347.7s 2.9×10-3s-1
1. Record your results in Table 12.1.
Make sure you adjust for your start times when you added HCl.
a. Each well contains eight drops of liquid. Remember that M1V1 = M2V2. Use this relation to calculate the molarity of Na2S2O3 in each well and record the values.
b. Record time to cloud to the point the X’s are fully hidden.
c. Assume that the same amount of sulfur precipitated in each well to hide the X’s. The reaction rate for each well is then precipitate/time, which is proportional to 1/t. Calculate and record 1/t.
2. Graph time (y-axis) vs. molarity (x-axis). Determine if the curve is linear.
3. Graph 1/time (y-axis) vs. molarity (x-axis). Determine if the curve is linear.
4. Determine the order of reaction.
Discussion and Conclusion:
Write your discussion and conclusion in paragraph form, using the following to guide your thoughts. Remember to support your statements with data from your investigation. If you use outside sources for your discussion, please cite them properly.
• Briefly summarize what you did.
• State the relation you observed between reaction rate and molarity of Na2S2O3 ?and the order of reaction.
• State the relation you observed between reaction rate and molarity of HCl and the order of reaction.
• Comment on how linear your data was in the time vs concentration and 1/ time vs. concentration graphs. Explain why you constructed both graphs and what information each provides.
• Explain how you determined the order of reaction.
• Comment on whether the reaction order appeared to be the same or not if you ?varied the concentration of HCl and not Na2S2O3.
• Were these results what you expected? Explain. Are there ways, other than ?varying concentration of the reactants, to determine order of reaction?
• Indicate sources of error in the investigation and suggest ways to avoid them in future experiments.
• Suggest improvements in the design of both parts of the experiment.
Discussion and Conclusion
For this laboratory experiment, we organized and realized an experiment with the objective of observing the relation between concentration and reaction rates, while determining the order of reactions and their mathematical implications, For this, we prepared test tubes with different ratios of 5 drops in total of Na2S2O3 and distilled water, followed by 3 drops of HCl in timed intervals of 15 seconds to measure and record the times they needed to cloud and their relation to concentration. After we finished the experiment and analyzed the resulting data, we observed that the reaction rate and molarity of Na2S2O2 were directly proportional; the higher the concentration of Na2S2O2 was, the faster the reaction occurred, reflecting in the reaction order as well. For the HCl however, we observed that it also showed a faster reaction rate when applied to test tubes with a higher molarity, and would therefore have a lower rate of reaction. For the analysis of the data we elaborated graphs to represent the data in a visual way and determine more quickly how each factor acted in relation to the other. In our first graph we compared time vs molarity and obtained a linear graph, demonstrating that they are directly proportional. For our second graph we used 1/time vs molarity and obtained a curved graph, which says that they are not necessarily proportional values.
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