Vitamin C Clock Reaction: Kinetics
Vitamin C Clock Reaction: Kinetics
Abstract
Collision theory refers to a theory that was proposed by Max Trautz and William Lewis independently. The theory qualitatively offers an explanation on the way chemical reactions take place and the reason behind differences in reaction rates in different reactions. The rate equation for a given chemical reaction refers to the equation that provides a link between the pressure or concentrations of the reactants and the parameters that are constant partial reaction orders and the rate coefficients. This experiment aimed to determine the rate of reaction of reaction between iodide and hydrogen peroxide and determine the order for the reactants and ultimately the rate constant. The time taken to turn all the reactants into product was measured using two trials with a number of runs that decreased in concentration of the reactant. Using the procedure for determining the initial rate, it is possible to determine the order for the reactants and ultimately the rate constant, k. From the experiment the rate of reaction, in trial 1, was determined to be 0.00217 Mol/L/s while in trial 2, the rate was 0.00268Mol/L/s. Using the rates and the concentrations for iodide and hydrogen peroxide, the rate constant in trial 1 was 0.282 Mol-1L-1s-1 while, in trial 2, the constant k was 0.2279 Mol-1L-1s-1. The differences in the rate constant may have been as a result of experimental errors.
Introduction
Collision theory refers to a theory that was suggested by Trautz and Lewis independently. The theory qualitatively offers an explanation on the way chemical reactions take place and the reason behind differences in reaction rates in different reactions. The theory argues that only a small portion of the collisions that take place when reactants hit one another may result in a significant chemical change. The collisions that are successful are known as successful or effective collisions. These collisions possess enough energy, also called the activation energy, during the collision that can break the bonds that exist and at the same time create new bonds. This leads in production of the products, in a reaction (Nic, Jirat, & Kosata, 2014).
The rate of a given reaction is usually dependent on the number of successful collisions that take place in the reaction. When more reactants collide there are high chances more products will be formed. When the concentration of the reactants is increased there are more particles that are available for collision and this increases the chances of having an effective collision. This enhances the rate at which the reaction occurs. Adding a catalyst in the reaction mixture reduces the energy that is needed for the chemical changes to occur. Consequently, the available energy will enable more particles to collide increasing the rate of reaction (Nic, Jirat, & Kosata, 2014).
In order to study the rates at which chemical reactions take place using chemical kinetics, the basic requirement for the reaction being studied is that the particles that are reacting will collide and have interaction with one another. The particles are able to collide since they possess some energy in them. The reacting particles also need to have enough energy that is needed for a successful collision that would lead to product production. Lack of this minimum energy will cause the reactant particles to bounce off each other after the collision with no reaction taking place. Other chemical reactions require that the particles that are reacting be in a specific orientation in order to provide a successful collision. When collisions take place in other different orientations, leads to failure in having an effective collision which leads to reduced in the rate of reaction (Godambe, 2001).
The rate equation for a given chemical reaction refers to the equation that provides a link between the pressure or concentrations of the reactants and the parameters that are constant partial reaction orders and the rate coefficients. In order to find out the rate equation for a givens system, the reaction rate and the mass balance for the system are combined. The rate of reaction may be determined by calculating the amount of a given product per second or by determining the speed at which the concentration of a given reactant in the reaction is reducing at any particular time. For instance, in a reaction where substances A and B are reacting, the rate of reaction may be determined by calculating the speed at which concentration of A is falling every second (Clark, 2002).
In this laboratory exercise, the rate of a chemical reaction was studied using a clock reaction. Clock reactions generate a rapid, sharp change in color when one of the reactants is entirely consumed. The clock reaction was thus used to determine the amount of time that was taken to have the reaction complete. The information obtained was used to determine the reaction rate and the reaction order where reactants fall. Using this procedure, the reaction rate for iodine that was reacting with hydrogen peroxide was determined.
The experiment involved two reactions, iodine and starch to give blue color and iodide and starch to give a colorless product. The time taken to turn all the reactants into product was measured using two trials with a number of runs that decreased in concentration of the reactant. Using the procedure for determining the initial rate, it is possible to determine the order for the reactants and ultimately the rate constant, k.
The experiment thus aimed to determine the rate of reaction of reaction between iodide and hydrogen peroxide and determine the order for the reactants and finally the rate constant.
Procedure
The first four reactants (0.25 M H2O2, 0.25 M NaI, 0.05 M Vitamin C, and 1 M acetic acid with starch) were added as shown in Table 1 below by measuring their volumes carefully I the labeled test tubes. The hydrogen peroxide was the last one to be added when timing was ready to be started. Timing was started when addition of hydrogen peroxide was started. When hydrogen peroxide was added, the content was stirred continuously, and the time taken for the reaction to complete recorded in Table 1 below.
Results
The volumes for the reactants and the time taken for the reaction to complete were recorded in Table 1 below. In trial 1, the time taken for the reaction to complete increased as the amount of iodine reduced while in trial 2 the time increased as the amount of hydrogen peroxide reduced.
Calculations
The actual concentration of H2O2 and I in each run were calculated as follows.
Concentration of iodide in trial 1 run 2 was
M=molesL
0.25M=moles0.0025 L=6.25×10-4moles of iodide
M=6.25×10-40.007 L=0.089 mol/L
Concentration of iodide in trial 1 run 3 was
M=molesL
0.25M=moles0.002 L=5.0×10-4moles of iodide
M=5.0×10-40.007 L=0.0714 mol/L
Concentration of iodide in trial 1 run 4 was
M=molesL
0.25M=moles0.0015 L=3.75×10-4moles of iodide
M=3.75×10-40.007 L=0.0536 mol/L
Concentration of iodide in trial 1 run 6 was
M=molesL
0.25M=moles0.0005 L=1.25×10-4moles of iodide
M=1.25×10-40.007 L=0.0179 mol/L
Concentration of hydrogen peroxide in trial 2 run 2 was
M=molesL
0.25M=moles0.0025 L=6.25×10-4moles of H2O2
M=6.25×10-40.007 L=0.089 mol/L
Concentration of hydrogen peroxide in trial 2 run 3 was
M=molesL
0.25M=moles0.002 L=5.0×10-4moles of H2O2
M=5.0×10-40.007 L=0.0714 mol/L
Concentration of hydrogen peroxide in trial 2 run 4 was
M=molesL
0.25M=moles0.0015 L=3.75×10-4moles of H2O2
M=3.75×10-40.007 L=0.0536 mol/L
Concentration of hydrogen peroxide in trial 1 run 6 was
M=molesL
0.25M=moles0.0005 L=1.25×10-4moles of H2O2
M=1.25×10-40.007 L=0.0179 mol/L
The rate of each run in concentration/time (s) was calculated as follows
Rate of reaction in trial 1 run 2 was
Rate of Reaction=Concentration (mol/L)Time in seconds
=0.089 mol/L41 s=2.17×10-3mol/l/s
Rate of reaction in trial 1 run 3 was
Rate of Reaction=Concentration (mol/L)Time in seconds
=0.0714mol/L49 s=1.46×10-3mol/l/s
Rate of reaction in trial 1 run 4 was
Rate of Reaction=Concentration (mol/L)Time in seconds
=0.0536 mol/L65 s=8.246×10-4mol/l/s
Rate of reaction in trial 1 run 6 was
Rate of Reaction=Concentration (mol/L)Time in seconds
=0.0179 mol/L345 s=5.188×10-5mol/l/s
Rate of reaction in trial 2 run 2 was
Rate of Reaction=Concentration (mol/L)Time in seconds
=0.089 mol/L 33s=2.69×10-3mol/l/s
Rate of reaction in trial 2 run 3 was
Rate of Reaction=Concentration (mol/L)Time in seconds
=0.0714 mol/L49 s=1.46×10-3mol/l/s
Rate of reaction in trial 2 run 4 was
Rate of Reaction=Concentration (mol/L)Time in seconds
=0.0536 mol/L74 s=7.24×10-4mol/l/s
Rate of reaction in trial 1 run 6 was
Rate of Reaction=Concentration (mol/L)Time in seconds
=0.0179 mol/L308 s=5.81×10-5mol/l/s
A graph of rate of reaction against concentration of iodide was plotted as shown in Figure 1 below
Figure 1: Rate of reaction against concentration of iodide
A graph of rate of reaction against concentration of iodide was plotted as shown in Figure 2 below
Figure 2: Rate of reaction against concentration of iodide
4) Determine what the graph would look like if the reaction order was 1 and if the reaction order was 2. Based on your determination what is the reaction order for both H2O2 and I- based on your data.
The order in which a reaction is can be determined by the nature of the curve of rate of reaction against the concentration. When the rate is directly proportional to the concentration, the reaction is a first order reaction, and the curve is a straight line with a positive gradient. When the rate of reaction is proportional to the power of concentration that is greater than one, the reaction is said to be in second order (Jim Clark, 2007). According to the graphs obtained both iodide and hydrogen peroxide reactions are first order reactions.
The rate constant (k) was calculated using run 2 for each trial as shown below.
For run 2 in trial 1, the rate constant was
r=k[H202][I-]
0.00217=k0.0890.107
=0.2279 /Mol/L/s
For run 2 in trial 2, the rate constant was
r=k[H202][I-]
0.00269=k0.0890.107
=0.282/Mol/L/s
The rate law for the reaction can thus be written as r=0.282[H202][I-] using the rate constant obtained in trial 1.
Discussion
The experiments involved two different reactions that took place simultaneously. In one reaction, two molecules of iodide reacted with hydrogen peroxide to give iodine and water. This was the reaction whose rate was being measured. The released iodine then reacted with vitamin C. The reaction between iodide and hydrogen peroxide is a slow process while the reaction between vitamin C and iodine occur at a fast rate. The elemental iodine produced in the first reaction gives a blue color when reacted with starch. In the presence of vitamin C, iodine reacts with vitamin C to give a colorless solution. After a short time during the reactions, the color remained clear as the produced iodine was being eliminated. It is at this point that Vitamin C creates a clock reaction which can last for 1 - 2 minutes. Once the Vitamin C is exhausted, the iodine element reacts with starch forming a blue-colored complex. Chemists determine the rate of reaction in by determining the appearance of a product or the disappearance of a reactant. In this reaction, the appearance of the product, iodine indicated by the appearance of the starch complex.
The rate constants that were calculated using the two trials gave values that were slightly different from one another. This difference was not expected and may have resulted from errors during the experiments. Some of the errors may have resulted during sampling, by measuring incorrect volumes or by timing the reaction time incorrectly. The graphs of rate of reaction against concentration of iodide and hydrogen peroxide indicate that there is a direct relationship between the concentration of the reactants and the speed of the reaction. As the concentration increases, there is also an increase in the rate of the reaction, and as the concentration decrease the rate of the reaction reduces.
The lab applied several principles of green chemistry. For instance, the experiment used safe reactants in the laboratory that did not create dangerous substances.
Conclusion
This experiment aimed to determine the rate of reaction of reaction between iodide and hydrogen peroxide and determine the order for the reactants and ultimately the rate constant. The experiment involved two reactions, iodine and starch to give blue color and iodide and starch to give a colorless product. The time taken to turn all the reactants into product was measured using two trials with a number of runs that decreased in concentration of the reactant. Using the procedure for determining the initial rate, it is possible to determine the order for the reactants and ultimately the rate constant, k. From the experiment, the rate of reaction in trial 1 was determined to be 0.00217 Mol/L/s while in trial 2, the rate was determined to be 0.00268Mol/L/s. Using the rates and the concentrations for iodide and hydrogen peroxide, the rate constant in trial 1 was 0.282 Mol-1L-1s-1 while, in trial 2, the constant k was 0.2279 Mol-1L-1s-1. The differences in the rate constant may have been as a result of experimental errors.
Reference List
Clark, J. (2002). The collision theory of reaction rates. Retrieved March 21, 2014, from http://www.chemguide.co.uk/physical/basicrates/introduction.html
Godambe, D. (2001). Kinetics and Collision Theory. Retrieved March 21, 2014, from http://www.harpercollege.edu/tm-ps/chm/100/dgodambe/thedisk/kinetic/6back.htm
Jim Clark. (2007). The Effect of Changing Conditions in Enzyme Catalysis. Retrieved March 22, 2014, from http://www.chemguide.co.uk/organicprops/aminoacids/enzymes2.html
Nic, M., Jirat, J., & Kosata, B. (2014). Collision theory. Retrieved March 21, 2014, from http://goldbook.iupac.org/C01170.html
