Overview
This Experiment seeks to introduce the general concepts of kinetics and of UV Vis spectroscopy as a branch of analytical chemistry.  In particular, the use of UV Vis spectroscopy and the isolation method in kinetics are explored. The experiment consists of two parts: I. Measurement of the concentration of chromium in solution using a UV-Vis spectrometer and II. Determination of kinetics for oxidation of iodide ion by hydrogen peroxide using the spectrometer
 
Safety
Good Chemical Practice must be followed throughout.
 
Consult a demonstrator or technician about the risk and hazards associated with chromium (III) nitrate before commencing the experiment.
 
Because of the low concentration of the solutions, no particular hazard is involved in the iodide oxidation.
 

  1. Measurement of the concentration of chromium in solution

Aim
To determine the concentration of Cr3+ in a solution of unknown composition.
 
Quantitative analysis of chromium nitrate solutions
To determine the concentration of any solute (something dissolved in a solvent) using the spectrophotometer, the first step is to run a spectrum of the solution (solute + solvent) using the solvent alone as the background.  In this way, the best wavelength to employ for any analytical measurements can be identified (typically the lmax of the most intense peak).  A series of solutions of known concentration of the solute are then made up and the absorbance of each solution at lmax measured and plotted as a function of concentration to produce a calibration curve.  This is then employed to calculate the extinction coefficient, e in M-1cm-1, at lmax.  The Molar Decadic Extinction Coefficient can then be employed to determine unknown concentrations by measuring the absorbance at lmax, or the concentration can be read off directly from the calibration curve.
 
Measurement of the UV-Vis spectrum and determination of lmax and the extinction coefficient
Prepare a set of solutions of chromium (III) nitrate (Cr(NO3)3.9H2O) of accurately known concentrations  in the range 0.010 – 0.10 M.  (Note the phrase “..accurately known concentrations ..”;  although the amount of chromium nitrate required to make a solution exactly 0.10 M is known, it is difficult to weigh exactly a predetermined amount.  Therefore an amount is accurately weighed that gives a concentration close to 0.1M.  Although the solution is not exactly the specified concentration, its actual concentration is accurately known).
 
(1)   Weigh accurately a quantity of chromium (III) nitrate to make 100 cm3 of approximately 0.10 M solution in water (the molecular weight is usually given on the label of the bottle).
(2)   Make the 0.10 M solution by transferring ALL the weighed mass of chromium (III) nitrate to a clean, 250 cm3 beaker, adding approximately 50 cm3 of distilled water, dissolving the solid, by placing a magnetic follower into the solution and placing the beaker onto a magnetic stirrer.  Transfer the solution to a 100 cm3 volumetric flask, rinsing the beaker and adding the rinsing’s to the flask, and finally adding water to the solution in the flask to make up to 100 cm3.
(3)   Into 5 separate 50 cm3 volumetric flasks, pipette 5, 10, l5, 20 and 25 cm3 of the 0.1M chromium (III) nitrate solution and make each flask up to the mark with distilled water.
 
 
A Demonstrator or Technician will instruct you in the use of the spectrophotometer.
 
Fill the carousel with the cuvettes (the cuvettes have a path length of 1cm) in the order shown in the diagram beside the spectrometer.  Record a spectrum from 350 – 700 nm for the solution in the cuvette in position 1 (0.1M Cr(NO3)3 solution), and employ this spectrum to identify lmax, the path length of the cuvette you will use is 1 cm.  Run the sample to give a spectrum of the solution, two peaks will be displayed, record the wavelength (λ /nm) and absorbance of both peaks. Run the sample a second time and record the wavelength (λ /nm) and absorbance of both peaks. Save the second run to a USB drive for inclusion in your report.
 
As you will see, the chromium nitrate spectrum has two main absorption bands; use the spectrum to determine which peak, and hence which lmax, you will use and justify your choice.  This should then be employed for the remainder of the experiment. The lmax value determined will be the fixed λ at  which the rest of the solutions will be measured at. The spectrometer should then be set to your chosen lmax (see a Demonstrator or Technician to change the settings on the spectrophotometer). The absorbance at lmax of all the solutions (except that in the cuvette in position 1, 0.1M) and including the unknown solution should be recorded.  Remember, lmax should not depend upon the concentration of the Cr species.  Summarize your data in an appropriate table.
 
Treatment of results
Plot a calibration curve of absorbance (Y-axis) against concentration (X-axis) for lmax, labelling the axes.  The plot should be linear through the origin, showing that the absorbance is directly proportional to the concentration of the chromium nitrate.  If your plot is not linear, or does not go through the origin, consult a demonstrator or technician.  Use your results to calculate the Molar Decadic Extinction Coefficient of chromium nitrate at lmax (don’t forget to include the units of e!).
 
Determination of an unknown concentration of Cr(NO3)3
A solution containing Cr(NO3)3 of unknown concentration is provided.  Run a spectrum of the solution (include in your report) and use it to determine the concentration of Cr(NO3)3.
 

  1. Measurement of the kinetics of iodide oxidation.

Aim
The aim of this experiment is to find how the rate of a reaction depends on the concentrations of the reactants. This is known as finding the rate law for the reaction. The reaction you will investigate is the oxidation of iodide by hydrogen peroxide:
 
H2O2 + 3I + 2H3O+ ↔ I3 + 4H2O                 (1)
 
Experimental procedure
Writing the rate law as:
Rate = k[H2O2]L[I]M[H3O+]N
-d[H2O2] = k[H2O2]L[I]M[H3O+]N     (2)
 
The procedure you will use involves measuring the concentration of I3 as it is formed by reaction (1).  I3 is determined spectrophotometrically. The absorbance is measured at the lmax of I3 at 350 nm, where I3 is the only absorbing species.   The concentration of I3 is calculated from absorbance according to:
[I3] = A/e350l   (3)
 
ε350 = 2.6 x 104 dm3 mol-1 cm-1.
 
De-ionised water must be used for all solutions.  Glassware should be rinsed well with de-ionised water before use.  Make sure all solutions are clearly labelled as soon as they are prepared. The reactions are carried out at room temperature.  (Kinetic measurements are usually carried out at a controlled temperature in a thermostat, but one is not employed in this experiment to keep the apparatus as simple as possible).
 
A stock solution of H2O2 of approximate concentration 0.02 M is provided.  Since H2O2 is unstable the concentration is not exact and must be determined during the experiment from the absorbance readings. Solutions of I (0.1M KI) and H2SO4 (1M) are also provided, these concentrations may be assumed to be accurate; they are also sufficiently larger than that of the peroxide that they can be considered to remain constant throughout the reaction as per the “isolation” approach.
 
Following the reaction to completion
Before starting the run, prepare a results table to record the time and the absorbance as the run is going on. Read and record the absorbance at 350 nm after a reaction time of 1 minute, and at 1 minute intervals for 20 minutes and then a final reading after 30 min.  Remember to set to zero absorbance with the reference cell (control solution) in position between absorbance readings.  Record the laboratory temperature at the start and end of the run.
 
Experimental procedure
Preparation of the dilute H2O2 solutions
The stock solution must be diluted by a factor of 200 before use and this can be done in one step. Thus, calculate the volume of the 0.02M H2O2 solution that should be diluted to 50 cm3 to give a concentration of 1 x 10-4 M (check with a Demonstrator or Technician your calculations are correct before preparing the solution).  Place this volume of H2O2 into a 50 cm3 volumetric flask and dilute to the mark with de-ionised water.
 
Prepare the spectrophotometer for measurement, set the wavelength to 350 nm prior to running the experiment.
 
A Demonstrator or Technician will instruct you in the use of the spectrophotometer:
 
Running the experiment
Use 50cm3 beakers (labelled) as reservoirs for the I, H2SO4, H2O and H2O2 solutions.
Using one pipette per solution, transfer the solutions, as required, to the reaction (R) and control
(C) beakers.
 
(1)   Place the control beaker onto a magnetic stirrer with a magnetic follower added. To the control beaker (C) add 10 cm3 of the I (KI) solution, 10 cm3 of the sulphuric acid (H2SO4) solution using glass pipettes.  Add 5cm3 of water using the digital pipette and then stir briefly. Fill a cuvette with the control solution and place into position 0 in the carousel.  The spectrophotometer is zeroed using the control solution prior to measuring the absorbance of the reaction solution. (This solution must be prepared as soon as possible before preparing the reaction solution).
(2)   Place the reaction beaker onto a magnetic stirrer with a magnetic follower added.  To the reaction (R) beaker add 10 cm3 of the I solution, 10 cm3 of H2SO4 using glass pipettes.  Place a fresh tip onto the digital pipette. Start the reaction by injecting rapidly 5 cm3 of the 1 x 10-4 M H2O2 solution from the digital pipette into the reaction mixture.  Turn the stirrer on and allow two rotations of the stirrer bar before removing the beaker. As rapidly as possible fill a cuvette with the reaction solution and place into position 1 in the carousel.
 
The first absorbance reading is t = 0 and when the first reading is taken start the stop clock for the rest of the experiment.
 
Treatment of results and discussion
(1).  Calculate the concentrations of I3 from the absorbance values (use equation (3)).  The final value (30 min), after all the H2O2 has reacted, is referred to as the infinity value of the product concentration. In this reaction, the infinity value of I3 is equal to the initial H2O2 concentration,

[H2O2]0 (stoichiometry is 1:1, see equation (1)). Calculate the concentration of H2O2 for each of the measured points using the measured I3 concentrations.  (Note: [H2O2] = [H2O2]0 – [I3]).
(2).  Plot a graph of [H2O2] (Y-axis) against reaction time (X-axis). The rate of this reaction at any time, t, during the reaction is, by definition, equal to -d[H2O2]/dt. This means, in terms of the experimental results, the rate is equal to the slope of the tangent to the curve you have just plotted.  The value of the rate is related to the reactant concentrations by equation (2), and we want to use the experimental results to evaluate the parameters, k, L, M and N in this equation.  Remember, the isolation method is being employed so the concentrations of I and H3O+ can be assumed to be constant, and equal to the initial values throughout the whole reaction.  (Check this statement by working out the percentage changes in each of the reactant concentrations, assuming the reaction has gone to completion).   As was stated in the Theory section of the manual, if the reaction is first order in H2O2, i.e. L = 1, equation (2) can be simplified to give equation (4):
 
-d[H2O2]/dt = kobs[H2O2]     (4)
 
where the observed rate constant kobs = [I]0M[H3O+]0N and [I]0 and [H3O+]0 are the initial, unchanging concentrations of iodide and hydroxonium ions, respectively.
 
To test if the reaction is first order in H2O2, you could measure the slope of the tangent to the curve and see if it is proportional to [H2O2], as predicted by equation (4).  However, it is difficult to do this accurately, and there is an easier way of analysing the results, which to use the integrated rate equation (see Theory section):
 
Loge[H2O2] = Loge[H2O2]0 – kobst      (5)
 
If the results fit this equation, a plot of Loge[H2O2] against t should be a straight line and the slope of the line is equal to -kobs and the intercept at t = 0 equal to loge[H2O2]0.
(3).  Calculate values for Loge[H2O2] at each measured point.  Plot a graph of Loge[H2O2] (Y-axis) against the reaction time, t (X-axis).
(4).  If you find that your plot is linear, you can conclude that the reaction is first order in H2O2.  In this case, calculate the value of the observed rate constant, kobs (with units), for these reaction conditions and  record  it,  together  with  the  values  of  [I]0,  [H+]0,  [H2O2]0  and the mean laboratory temperature during the run.  If you are not convinced that your plot is linear, discuss your results with a demonstrator.
 
Questions 

  1. Explain why it is preferable to work with absorbance values in the range above about 0.1 and below about 1.0. /4
  2. Is the molar absorption coefficient (extinction coefficient) constant for the different concentrations of chromic nitrate solutions or should it be a variable? Explain your reasoning.

/2

  1. Explain briefly how you would adapt the experiment to determine the values of M and N in equation (2). /4
  2. Describe briefly the process which typically occurs in a molecule when it absorbs UV/VIS radiation. /3
  3. Why is it important the cuvette is free from bubbles, smudges and scratches? (Be specific!).

                                                                                                                                                   /2
 
Mark scheme

  Marks
Notebook /10
General aspects /30
Presentation /5
 
Structure /5
References /5
Added Value /5
Description of experiments /5
Discussion of data /5
Specific aspects /45
Correct identification of two lmax /2
Calibration curve correctly plotted and labelled /10
Justification for choice of lmax /2
Calculation of unknown concentration /5
Graph of H2O2 correctly plotted and labelled /10
Calculation of % change in concentration assuming 100% conversion /2
Graph of Loge[H2O2] vs t correctly plotted and labelled /10
Order with respect to H2O2 identified, kobs calculated /4
Questions /15
TOTAL /100

 
 
Experiment C1: Ultraviolet-visible Spectroscopy of Inorganic Solutions and its Application in Kinetic Measurements
 
Overview
This Experiment seeks to introduce the general concepts of kinetics and of UV Vis spectroscopy as a branch of analytical chemistry.  In particular, the use of UV Vis spectroscopy and the isolation method in kinetics are explored. The experiment consists of two parts: I. Measurement of the concentration of chromium in solution using a UV-Vis spectrometer and II. Determination of kinetics for oxidation of iodide ion by hydrogen peroxide using the spectrometer
 
Safety
Good Chemical Practice must be followed throughout.
 
Consult a demonstrator or technician about the risk and hazards associated with chromium (III) nitrate before commencing the experiment.
 
Because of the low concentration of the solutions, no particular hazard is involved in the iodide oxidation.
 

  1. Measurement of the concentration of chromium in solution

Aim
To determine the concentration of Cr3+ in a solution of unknown composition.
 
Quantitative analysis of chromium nitrate solutions
To determine the concentration of any solute (something dissolved in a solvent) using the spectrophotometer, the first step is to run a spectrum of the solution (solute + solvent) using the solvent alone as the background.  In this way, the best wavelength to employ for any analytical measurements can be identified (typically the lmax of the most intense peak).  A series of solutions of known concentration of the solute are then made up and the absorbance of each solution at lmax measured and plotted as a function of concentration to produce a calibration curve.  This is then employed to calculate the extinction coefficient, e in M-1cm-1, at lmax.  The Molar Decadic Extinction Coefficient can then be employed to determine unknown concentrations by measuring the absorbance at lmax, or the concentration can be read off directly from the calibration curve.
 
Measurement of the UV-Vis spectrum and determination of lmax and the extinction coefficient
Prepare a set of solutions of chromium (III) nitrate (Cr(NO3)3.9H2O) of accurately known concentrations  in the range 0.010 – 0.10 M.  (Note the phrase “..accurately known concentrations ..”;  although the amount of chromium nitrate required to make a solution exactly 0.10 M is known, it is difficult to weigh exactly a predetermined amount.  Therefore an amount is accurately weighed that gives a concentration close to 0.1M.  Although the solution is not exactly the specified concentration, its actual concentration is accurately known).
 
(1)   Weigh accurately a quantity of chromium (III) nitrate to make 100 cm3 of approximately 0.10 M solution in water (the molecular weight is usually given on the label of the bottle).
(2)   Make the 0.10 M solution by transferring ALL the weighed mass of chromium (III) nitrate to a clean, 250 cm3 beaker, adding approximately 50 cm3 of distilled water, dissolving the solid, by placing a magnetic follower into the solution and placing the beaker onto a magnetic stirrer.  Transfer the solution to a 100 cm3 volumetric flask, rinsing the beaker and adding the rinsing’s to the flask, and finally adding water to the solution in the flask to make up to 100 cm3.
(3)   Into 5 separate 50 cm3 volumetric flasks, pipette 5, 10, l5, 20 and 25 cm3 of the 0.1M chromium (III) nitrate solution and make each flask up to the mark with distilled water.
 
 
A Demonstrator or Technician will instruct you in the use of the spectrophotometer.
 
Fill the carousel with the cuvettes (the cuvettes have a path length of 1cm) in the order shown in the diagram beside the spectrometer.  Record a spectrum from 350 – 700 nm for the solution in the cuvette in position 1 (0.1M Cr(NO3)3 solution), and employ this spectrum to identify lmax, the path length of the cuvette you will use is 1 cm.  Run the sample to give a spectrum of the solution, two peaks will be displayed, record the wavelength (λ /nm) and absorbance of both peaks. Run the sample a second time and record the wavelength (λ /nm) and absorbance of both peaks. Save the second run to a USB drive for inclusion in your report.
 
As you will see, the chromium nitrate spectrum has two main absorption bands; use the spectrum to determine which peak, and hence which lmax, you will use and justify your choice.  This should then be employed for the remainder of the experiment. The lmax value determined will be the fixed λ at  which the rest of the solutions will be measured at. The spectrometer should then be set to your chosen lmax (see a Demonstrator or Technician to change the settings on the spectrophotometer). The absorbance at lmax of all the solutions (except that in the cuvette in position 1, 0.1M) and including the unknown solution should be recorded.  Remember, lmax should not depend upon the concentration of the Cr species.  Summarize your data in an appropriate table.
 
Treatment of results
Plot a calibration curve of absorbance (Y-axis) against concentration (X-axis) for lmax, labelling the axes.  The plot should be linear through the origin, showing that the absorbance is directly proportional to the concentration of the chromium nitrate.  If your plot is not linear, or does not go through the origin, consult a demonstrator or technician.  Use your results to calculate the Molar Decadic Extinction Coefficient of chromium nitrate at lmax (don’t forget to include the units of e!).
 
Determination of an unknown concentration of Cr(NO3)3
A solution containing Cr(NO3)3 of unknown concentration is provided.  Run a spectrum of the solution (include in your report) and use it to determine the concentration of Cr(NO3)3.
 

  1. Measurement of the kinetics of iodide oxidation.

Aim
The aim of this experiment is to find how the rate of a reaction depends on the concentrations of the reactants. This is known as finding the rate law for the reaction. The reaction you will investigate is the oxidation of iodide by hydrogen peroxide:
 
H2O2 + 3I + 2H3O+ ↔ I3 + 4H2O                 (1)
 
Experimental procedure
Writing the rate law as:
Rate = k[H2O2]L[I]M[H3O+]N
-d[H2O2] = k[H2O2]L[I]M[H3O+]N     (2)
 
The procedure you will use involves measuring the concentration of I3 as it is formed by reaction (1).  I3 is determined spectrophotometrically. The absorbance is measured at the lmax of I3 at 350 nm, where I3 is the only absorbing species.   The concentration of I3 is calculated from absorbance according to:
[I3] = A/e350l   (3)
 
ε350 = 2.6 x 104 dm3 mol-1 cm-1.
 
De-ionised water must be used for all solutions.  Glassware should be rinsed well with de-ionised water before use.  Make sure all solutions are clearly labelled as soon as they are prepared. The reactions are carried out at room temperature.  (Kinetic measurements are usually carried out at a controlled temperature in a thermostat, but one is not employed in this experiment to keep the apparatus as simple as possible).
 
A stock solution of H2O2 of approximate concentration 0.02 M is provided.  Since H2O2 is unstable the concentration is not exact and must be determined during the experiment from the absorbance readings. Solutions of I (0.1M KI) and H2SO4 (1M) are also provided, these concentrations may be assumed to be accurate; they are also sufficiently larger than that of the peroxide that they can be considered to remain constant throughout the reaction as per the “isolation” approach.
 
Following the reaction to completion
Before starting the run, prepare a results table to record the time and the absorbance as the run is going on. Read and record the absorbance at 350 nm after a reaction time of 1 minute, and at 1 minute intervals for 20 minutes and then a final reading after 30 min.  Remember to set to zero absorbance with the reference cell (control solution) in position between absorbance readings.  Record the laboratory temperature at the start and end of the run.
 
Experimental procedure
Preparation of the dilute H2O2 solutions
The stock solution must be diluted by a factor of 200 before use and this can be done in one step. Thus, calculate the volume of the 0.02M H2O2 solution that should be diluted to 50 cm3 to give a concentration of 1 x 10-4 M (check with a Demonstrator or Technician your calculations are correct before preparing the solution).  Place this volume of H2O2 into a 50 cm3 volumetric flask and dilute to the mark with de-ionised water.
 
Prepare the spectrophotometer for measurement, set the wavelength to 350 nm prior to running the experiment.
 
A Demonstrator or Technician will instruct you in the use of the spectrophotometer:
 
Running the experiment
Use 50cm3 beakers (labelled) as reservoirs for the I, H2SO4, H2O and H2O2 solutions.
Using one pipette per solution, transfer the solutions, as required, to the reaction (R) and control
(C) beakers.
 
(1)   Place the control beaker onto a magnetic stirrer with a magnetic follower added. To the control beaker (C) add 10 cm3 of the I (KI) solution, 10 cm3 of the sulphuric acid (H2SO4) solution using glass pipettes.  Add 5cm3 of water using the digital pipette and then stir briefly. Fill a cuvette with the control solution and place into position 0 in the carousel.  The spectrophotometer is zeroed using the control solution prior to measuring the absorbance of the reaction solution. (This solution must be prepared as soon as possible before preparing the reaction solution).
(2)   Place the reaction beaker onto a magnetic stirrer with a magnetic follower added.  To the reaction (R) beaker add 10 cm3 of the I solution, 10 cm3 of H2SO4 using glass pipettes.  Place a fresh tip onto the digital pipette. Start the reaction by injecting rapidly 5 cm3 of the 1 x 10-4 M H2O2 solution from the digital pipette into the reaction mixture.  Turn the stirrer on and allow two rotations of the stirrer bar before removing the beaker. As rapidly as possible fill a cuvette with the reaction solution and place into position 1 in the carousel.
 
The first absorbance reading is t = 0 and when the first reading is taken start the stop clock for the rest of the experiment.
 
Treatment of results and discussion
(1).  Calculate the concentrations of I3 from the absorbance values (use equation (3)).  The final value (30 min), after all the H2O2 has reacted, is referred to as the infinity value of the product concentration. In this reaction, the infinity value of I3 is equal to the initial H2O2 concentration,

[H2O2]0 (stoichiometry is 1:1, see equation (1)). Calculate the concentration of H2O2 for each of the measured points using the measured I3 concentrations.  (Note: [H2O2] = [H2O2]0 – [I3]).
(2).  Plot a graph of [H2O2] (Y-axis) against reaction time (X-axis). The rate of this reaction at any time, t, during the reaction is, by definition, equal to -d[H2O2]/dt. This means, in terms of the experimental results, the rate is equal to the slope of the tangent to the curve you have just plotted.  The value of the rate is related to the reactant concentrations by equation (2), and we want to use the experimental results to evaluate the parameters, k, L, M and N in this equation.  Remember, the isolation method is being employed so the concentrations of I and H3O+ can be assumed to be constant, and equal to the initial values throughout the whole reaction.  (Check this statement by working out the percentage changes in each of the reactant concentrations, assuming the reaction has gone to completion).   As was stated in the Theory section of the manual, if the reaction is first order in H2O2, i.e. L = 1, equation (2) can be simplified to give equation (4):
 
-d[H2O2]/dt = kobs[H2O2]     (4)
 
where the observed rate constant kobs = [I]0M[H3O+]0N and [I]0 and [H3O+]0 are the initial, unchanging concentrations of iodide and hydroxonium ions, respectively.
 
To test if the reaction is first order in H2O2, you could measure the slope of the tangent to the curve and see if it is proportional to [H2O2], as predicted by equation (4).  However, it is difficult to do this accurately, and there is an easier way of analysing the results, which to use the integrated rate equation (see Theory section):
 
Loge[H2O2] = Loge[H2O2]0 – kobst      (5)
 
If the results fit this equation, a plot of Loge[H2O2] against t should be a straight line and the slope of the line is equal to -kobs and the intercept at t = 0 equal to loge[H2O2]0.
(3).  Calculate values for Loge[H2O2] at each measured point.  Plot a graph of Loge[H2O2] (Y-axis) against the reaction time, t (X-axis).
(4).  If you find that your plot is linear, you can conclude that the reaction is first order in H2O2.  In this case, calculate the value of the observed rate constant, kobs (with units), for these reaction conditions and  record  it,  together  with  the  values  of  [I]0,  [H+]0,  [H2O2]0  and the mean laboratory temperature during the run.  If you are not convinced that your plot is linear, discuss your results with a demonstrator.
 
Questions 

  1. Explain why it is preferable to work with absorbance values in the range above about 0.1 and below about 1.0. /4
  2. Is the molar absorption coefficient (extinction coefficient) constant for the different concentrations of chromic nitrate solutions or should it be a variable? Explain your reasoning.

/2

  1. Explain briefly how you would adapt the experiment to determine the values of M and N in equation (2). /4
  2. Describe briefly the process which typically occurs in a molecule when it absorbs UV/VIS radiation. /3
  3. Why is it important the cuvette is free from bubbles, smudges and scratches? (Be specific!).

                                                                                                                                                   /2
 
Mark scheme

  Marks
Notebook /10
General aspects /30
Presentation /5
 
Structure /5
References /5
Added Value /5
Description of experiments /5
Discussion of data /5
Specific aspects /45
Correct identification of two lmax /2
Calibration curve correctly plotted and labelled /10
Justification for choice of lmax /2
Calculation of unknown concentration /5
Graph of H2O2 correctly plotted and labelled /10
Calculation of % change in concentration assuming 100% conversion /2
Graph of Loge[H2O2] vs t correctly plotted and labelled /10
Order with respect to H2O2 identified, kobs calculated /4
Questions /15
TOTAL /100