I need help answering questions 1-7. I have provided all the data needed

Introduction. Kinetics is the study of the rates of chemical reactions. In this experiment you will be studying the rate of the redox reaction between the arsenic acid and the iodide ion in aqueous solution. You will investigate how the rate changes when you vary the concentration of the reactants and also when you vary the temperature. You will use these experimental measurements to determine the rate law and to calculate the activation energy for the reaction. You can read about the basic ideas of kinetics- including rate laws, rate constants, mechanisms, and activation energy- in the Kinetics chapter of your textbook.

The iodine clock reaction refers to reactions that change color from colorless to a blue color due to the formation of the triiodide ion combining with starch. This change in color come in very exact timing.

The iodine clock reaction that you will be studying is the following:

6 H+(aq)  + IO3-(aq)  + 8 I-(aq)  -> 3 I3-(aq) + 3 H2O(l) (RXN #1)

To better be able to study the timing of this reaction we add a reducing agent, arsenious acid which reacts rapidly with the triiodide ion as shown in RXN 2 below, and thus delaying the blue color. Once all the arsenious acid is reacted then the color changes to blue due to the triiodide-starch complex.

H3AsO3(aq)  + I3-(aq)  + H2O(l)   HAsO42-(aq)  + 3 I-(aq)  + 4 H+(aq) (RXN #2)

Here's how the reaction timing works. The initial reaction mixture will contain relatively large amounts of iodate and iodide; a small, precisely known amount of arsenious acid; plus a few drops of starch indicator (which will turn blue when complexed with the triiodide ion). When all the reagents are mixed together at time zero, reaction #1 will begin and I3- will be created as a product. But all of the I3- created will quickly be consumed in reaction #2 and it will not have time to interact with the starch. This will be the situation until all the arsenious acid (H3AsO3(aq)) is used up. Once the arsenious acid is gone, the I3- created in reaction #1 will begin to accumulate in the solution and the triiodide and starch will stick together to make a blue-colored complex. So, when your reaction solution turns blue, you are really measuring the time it takes for all the arsenious acid to be consumed. Because you know precisely the moles of arsenious acid you started with, you should be able to use the stoichiometric relationship between arsenious acid and the triiodide ion to calculate how many moles of iodate ion were consumed in that same interval of time, which will allow you to calculate the initial rate of the reaction.

You will then repeat the experiment using different initial concentrations of iodate and iodide ions to see how that affects the rate of the reaction. This information will allow you to determine the experimental rate law for the reaction. Finally, you will repeat the experiment at different temperatures to see how that affects the rate of reaction and the rate constant, k. By plotting a graph of ln k as a function of 1/T, you will be able to calculate the activation energy for the reaction.

You will access the virtual lab at the following link: http://web.mst.edu/~gbert/IClock/Clock.html

The first step is to choose Prepare Solutions. You will pick concentrations for reactants in solution A and B and temperatures with drop down menus. The chosen concentrations represent the stock concentrations. The initial concentrations are on the right. Make sure to note both stock and initial concentrations in your data tables in your notebook.

The following are your choices:

Solution A : 50mL

Solution A will contain KIO3, H+ (pH=-log[H+]), H3AsO3(aq), and starch.

The stock [KIO3]can be 0.010 M, 0.020M, 0.040M.

The Buffer pH can be 4.40, 4.70, 5.00.

The [H3AsO3] is 0.003M.

The starch is 0.006%.

Solution B : 50mL

Solution B will contain KI.

The [KI] can be 0.10M, 0.20M, 0.30M.

Temperature

The reaction temperature can be 5C, 15C, 25C, 35C, 45C.

Choose Mix the solution, and then you should see an image of two Erlenmeyer flasks and a mixer in between. Above this image, you will see the timer settings: start, stop timer, ReSet.

Choose Start- the solution will be mixed, and the timer will begin.

Choose Stop when the solution has the first appearance of blue.

Record this time.

ReSet allows for the repeat of the trial with the same concentrations and temperature.

To change the concentrations for a new trial, chose prepare solutions (return to step 1).

You should run enough trials to determine the reaction orders of IO3-, H+, and I-. Make sure to run each trial in duplicate and then use average of times measured in calculations.

You will follow the same steps but, in these experiments, you will vary the reaction temperature while keeping the initial concentrations constant. Make sure to do all available temperatures (5C, 15C, 25C, 35C, 45C) and again run each temperature at least twice. Again, record all of your data in your notebook.

Now that you have finished the experimental part of the experiment, you should have all the following recorded in your lab notebook, hopefully in a neat, organized way:

• The composition of each reaction mixture (solution A and B).

• The time and temperature for each of the experimental mixtures. Note any additional observations.
• At the end of the experiment, put all of your time and temperature data into a neat, organized table. You should include the actual values recorded in lab, as well as the average values you will be using in your calculations.

The lab write-up (data analysis) will be done by completing all the steps on pages 0-5 to 0-10 (there are hints for how to do these calculations below). The “lab report” that you are required to hand in consists of three things:

1. Your lab notebook pages

2. The data analysis on pages

3. A printed copy of your computer graph of ln k versus 1/T.

These are the only things you need to hand in. You are not required to hand in a typed, formal written lab report for this experiment. Remember, this lab report is individual work- if your write-up or graph looks identical to any other student’s, I will assume it has been copied.

Calculating the Reaction Rate. Remember, the starch turns blue when all the arsenious acid has been consumed. To calculate the reaction rate (the initial rate) for each trial:

1. The initial concertation of arsenious acid is completely consumed in reaction #2 in the measured time interval.

The formula is:

2. The rate of arsenious acid consumption can be related to the rate of iodate ion consumed and the rate of Reaction #1 by using the stoichiometric relationships.

3. Use initial concentrations to determine orders of reaction relative to reactants.

Note: In this calculation we are assuming that the rate stays essentially constant during the measured time. This is a reasonable assumption since we are only running the reaction for a short period of time, so the [IO31-] and the [I–] don’t change very much compared to their initial values.

Calculating the Activation Energy. The relationship between the rate constant, the temperature, and the activation energy is given by the Arrhenius equation:

Where k is the rate constant, Ea is the activation energy, R is the ideal gas constant, T is the absolute temperature, and A is the Arrhenius constant (or frequency factor).

If you take the natural log of both sides of this equation, you can derive the relationship, which can be written as

If you examine this relationship you will see that it is a linear equation (y = mx + b), where y = ln k, x = 1/T, and the slope is equal to -Ea/R. So to experimentally determine the activation energy for a reaction, you run the reaction at different temperatures (all other factors held constant) and then make a graph of ln k as a function of 1/T. The slope of this line is -Ea/R.

In your experiment, identical reagent concentrations were used, while the temperature of the reaction was varied Once you calculate the values of the rate constants (k) for these 5 different temperatures. Use Excel to make a graph in which the independent variable (x) is 1/T (be sure to use temperature in Kelvin) and the dependent variable (y) is ln(k). It is usually best to use the "scatter" mode when you make the graph, then add a trend line (using linear regression) to fit the best straight line to the data. Have Excel print the equation of this line right and the correlation coefficient on your graph so you will have the value of the slope. With the value of the slope you can calculate the activation energy using the relationship: slope = –Ea/R.

1. Reaction Rate. Calculate the initial reaction rates (in M/s) for each of your mixtures. Show your calculations below. (Since you did more than one trial, you can do one calculation using the average time value for your two trials as long as the two temperatures are the same.) Make sure to label your numbers with the correct units.

2. Determining the Rate Law. By considering how the initial rate changed when you changed the initial concentration of either iodate or iodide or H+, write the experimental rate law for the reaction below. Show all work in determining the reaction orders.

What is the overall order of the reaction? _________

3. Calculating k values. Using your calculated initial reaction rates and the experimental rate law, calculate the value of k, the rate constant, for each of the trial mixtures include units. Be sure to label your numbers with units. Show your work below. Final result should be an average k ( 2 x standard deviation).

4. Fill out the summary table below or Make table on Excel and print out. The table includes columns for the natural log of k and 1/T because these values will be used for your graph. Make sure you keep at least 3 significant figures for the k, ln(k), and 1/T values- if you don't you may get rounding-off errors that will affect your graph.

Summary Table of Calculated Values

*since you did multiple trials, please only include average values in this table

5. Finding Activation Energy. Make a graph of ln k as a function of 1/T using the data from mixtures with varied temperatures (which all had the same reagent composition). Be sure the graph has a specific title and properly labeled axes. When making the graph, use the scatter mode and then add a linear regression trend line- be sure to have Excel print the equation for the line, and the R2 value on the graph. Attach a copy of your graph, which should be done using to the report.

Using the slope of your graph, as determined by linear regression, calculate the value of the activation energy (in kJ/mole) for the reaction. Be sure to label your numbers with units. Show this calculation below:

6. What is your R2 value for your linear regression line? Comment on this value and what it tells you about your data.

7. Sketch a potential energy diagram for this reaction on the axes below. For this reaction, ∆H = –300. kJ/mole. Use your experimental value for the activation energy. Each square on the energy axis represents 50 kJ/mole. Be sure to label the reactants, the products, the transition state, ∆H and Ea on your diagram.