Reaction Kinetics<\/p>\n
The rate of a chemical reaction can be expressed in a variety of ways. We can express rate in terms of the disappearance of reactant or the formation of product over a given amount of time. Often the rate expression we choose to determine experimentally depends on the feasiblity of detection. Because usually there will be one molecule in our reaction that is easier to detect than the others. In this activity you will observe how reactions involving 2 purple substances, iodine and crystal violet, can be detected in solution to observe and study reaction rates. <\/p>\n
PART I: Deriving the rate law of the Iodine Clock Reaction<\/p>\n
There are many versions of the iodine clock reaction. A common theme of these clock reactions is oxidation of iodide (I-) to elemental iodine (I2). In this lab we will observe the Landolt clock reaction: <\/p>\n
IO3\u2212\u00a0(aq) + 3 HSO3\u2212\u00a0(aq) \u2192 I\u2212\u00a0(aq) + 3 HSO4- (aq)<\/p>\n
IO3\u2212\u00a0(aq) + 5 I\u2212\u00a0(aq) + 6 H+\u00a0(aq) \u2192 3 I2\u00a0(aq) + 3 H2O (aq)<\/p>\n
I2\u00a0(aq) + HSO3\u2212\u00a0(aq) + H2O (aq) \u2192 2 I\u2212\u00a0(aq) + HSO4\u2212\u00a0(aq) + 2 H+ (aq)<\/p>\n
To determine the rate law for this reaction, an experiment must be performed to determine the order with respect to each reactant. The first step is much slower than steps 2 and 3, thus step 1 is called the \u201crate-limiting step.\u201d As I2 forms in the 2nd step it is immediately consumed by the excess bisulfite as shown in the 3rd step. When the HSO3- is completely consumed, I2 will build up and precipitate out of solution. However, if starch is added to the solution prior to the reaction, the I2 will complex with the starch turning the solution a deep purple color. The iodine starch reaction proceeds as follows:<\/p>\n
I2 (aq) + starch (aq) I2 starch complex (aq) (deep blue\/purple) <\/p>\n
Because of this color change, the rate of reaction can be observed visually. We need to record the time from when the solutions are first mixed until the appearance of this deep purple color. At this point we can say that all of the HSO3- has been consumed thus [I2] = [HSO3-]. So it will be easiest to express reaction rate in terms of the concentration of the reactant HSO3-. <\/p>\n
\uf044[HSO3-] \/ \uf044t = rate = k [IO3-]x [HSO3-]y [H+]z<\/p>\n
Rate laws must be derived empirically to determine the values of x, y and z. One method is to run this reaction 3 times. The first reaction will be referred to as the control. The rate of this reaction will be compared to the other 2. The second reaction will involve doubling [IO3-] while holding [HSO3-] constant relative to the first reaction. The third reaction will involve doubling [HSO3-] while holding [IO3-] constant relative to the first reaction. You will also observe a fourth reaction for the purpose of evaluating the effect of H+ concentration on reaction rate. <\/p>\n
Complete the PROCEDURE and RESULTS\/DISUSSION sections for part I below.<\/p>\n
PROCEDURE: Use this video from Mr GrodskyChemistry to write a simple procedure for this lab. <\/p>\n
Some things to note:<\/p>\n
The chemist is mixing solution 1 with solution 2 before starting the timer. <\/p>\n
Solution 1: KIO3, H2O and Starch <\/p>\n
Solution 2: NaHSO3 and H2O<\/p>\n
50 mL of each solution are mixed in a 100 mL beaker in each reaction.<\/p>\n
The video shows a table with the concentrations of each reaction in the final reaction mixture. Record this table in the Results section below. In your table, record the approximate reaction times for each of the 4 reactions.<\/p>\n
In your procedure include waste disposal guidelines. Include information on the relative toxicity of the chemicals used in this lab. Can anything be disposed of in the common drain? One place to find this information is in the Material Safety Data Sheet (MSDS) for each chemical. The MSDS for these chemicals should be readily available and easily searchable online.<\/p>\n
Add your procedure below<\/p>\n
RESULTS\/DISCUSSION:<\/p>\n
Include your data table here.<\/p>\n
Calculate the rate of each reaction by considering that the solution only turns purple\/blue when all of the HSO3- is consumed. Thus, if you know the concentration change of HSO3- (given in the table) and you know the amount of time it took for the solution to turn blue, then you can calculate \uf044HSO3- \/ \uf044t.<\/p>\n
Compare the rates of reactions 2, 3 and 4 to reaction 1. Note that if the rate doubles (2x) when the concentration of a reactant doubles then the reaction is 1st order with respect to that reactant. If the rate increases by a factor of 4 (22) when the concentration of a reactant is doubled then the reaction is 2nd order with respect to that reactant and if the rate increases by a factor of 8 (23) when the reactant concentration doubles then the reaction is 3rd order with respect to that reactant.<\/p>\n
Once you have used the method above to determine x, y and z then use any single point in your table to input values of concentration and rate and solve for k. Solve for k for each of the 4 points and report the average k.<\/p>\n
What to report: <\/p>\n
The average rate constant for the 4 reactions<\/p>\n
Show your work in solving for k in each reaction.<\/p>\n
The overall rate law for this reaction.<\/p>\n
Describe your comfort level with deriving rate laws using these experimental techniques.<\/p>\n
Would it have been helpful to perform this lab yourself in person? Or were you able to learn this by watching videos? Explain.<\/p>\n
PART II: Monitoring reaction of Crystal Violet to derive an integrated rate law. <\/p>\n
Integrated rate laws allow us to calculate the reactant concentration at any given time in a reaction. Of course the rate of disappearance of reactant will depend on the reactant concentration itself. So the initial rate of disappearance will not equal the rate at any later time in the reaction. By monitoring the reactant concentration over time we can observe how rate is affected by concentration which allows us to determine if the rate law for the reaction is 0th, 1st or 2nd order. In this exercise you will use data to derive the integrated rate law for the following reaction:<\/p>\n
CV+ (aq) + OH \u2013(aq) CVOH (1)<\/p>\n
(violet\/purple) colorless<\/p>\n
In this lab you will measure the concentration of CV+ as a function of time. CV+ changes from a purple color to a clear\/colorless solution so it can be monitored spectrophotometrically. The Spectrophotometer shines specific wavelengths of light onto the sample and measures the fraction of that light that passes through the sample as a measure of \u201ctransmittance.\u201d The amount of light that passes through the sample is dependent on the molar absorptivity of the molecules and the concentration (Beer\u2019s Law). As the concentration of CV+ decreases during the reaction, we expect the absorbance to decrease also. <\/p>\n
Complete the PROCEDURE and DATA ANALYSIS section for part II below.<\/p>\n
PROCEDURE: Use this video from The University of Texas El Paso to write a basic procedure for this lab. <\/p>\n
The analyst mixes the crystal violet and sodium hydroxide solutions then quickly pours them into the cuvette. Why is it important to do this as quickly as possible?<\/p>\n
How do you expect the color to change over the course of the reaction? What will the color be at the beginning and end of the reaction?<\/p>\n
What concentration of CV is used in this experiment? What are the 2 concentrations of NaOH used?<\/p>\n
Why does the analyst use water to calibrate the spectrophotometer (~2:30 and 3:50)? What percentage of light is expected to be absorbed by the water? What percentage of light will be transmitted through the water? What does the word \u201ccalibration\u201d mean in this context?<\/p>\n
In this experiment, the analysts determine k’ instead of k. What is k’ and why is it useful to consider in place of k?<\/p>\n
DATA ANALYSIS<\/p>\n
Use the sample data give for OL7-Reaction Kinetics.xls to prepare a plot of Abs vs time, ln Abs vs time and 1\/Abs vs time to determine if the reaction of is 0th, 1st or 2nd order with respect to CV. Note: The absorbance is proportional to but not equal to concentration. In this lab we just need to monitor the change in absorbance with time and this change would be the same if we could measure the change in concentration with time.<\/p>\n","protected":false},"excerpt":{"rendered":"
Reaction Kinetics The rate of a chemical reaction can be expressed in a variety of ways. We can express rate in terms of the disappearance of reactant or the formation of product over a given amount of time. Often the rate expression we choose to determine experimentally depends on the feasiblity of detection. Because usually […]<\/p>\n","protected":false},"author":1,"featured_media":0,"comment_status":"closed","ping_status":"closed","sticky":false,"template":"","format":"standard","meta":{"footnotes":""},"categories":[1],"tags":[10],"class_list":["post-79048","post","type-post","status-publish","format-standard","hentry","category-research-paper-writing","tag-writing"],"_links":{"self":[{"href":"https:\/\/papersspot.com\/blog\/wp-json\/wp\/v2\/posts\/79048","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/papersspot.com\/blog\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/papersspot.com\/blog\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/papersspot.com\/blog\/wp-json\/wp\/v2\/users\/1"}],"replies":[{"embeddable":true,"href":"https:\/\/papersspot.com\/blog\/wp-json\/wp\/v2\/comments?post=79048"}],"version-history":[{"count":0,"href":"https:\/\/papersspot.com\/blog\/wp-json\/wp\/v2\/posts\/79048\/revisions"}],"wp:attachment":[{"href":"https:\/\/papersspot.com\/blog\/wp-json\/wp\/v2\/media?parent=79048"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/papersspot.com\/blog\/wp-json\/wp\/v2\/categories?post=79048"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/papersspot.com\/blog\/wp-json\/wp\/v2\/tags?post=79048"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}