Laboratory Exercise 2: Chemical Kinetics Cleans Up
Aim:
In this week’s lab you will:
Monitor changes in absorbance of a reactant using UV/Vis spectroscopy.
Determine the rate law & rate constant for the reaction between a dye, such as, Fast Green FCF and sodium hypochlorite (NaOCl, household bleach).
Introduction
Chemistry is concerned with changes such as those that govern how reactants are converted to products. Determining how fast a reaction will occur is critical for many commercial processes as well as many metabolic reactions in the cell. In addition, factors such as concentration of substrates, temperature, physical states of reactants and presence of catalysts all affect the rate of a reaction.
In this laboratory we will be determining the effect of concentration on reaction rates – how quickly household bleach reacts with a dye. This is easily observed by adding the two solutions and watching to see how quickly the solution becomes clear.
A rate law is a mathematical expression illustrating how the rate of a reaction depends on, for example, concentration and involves a rate constant. Eg. For the generic reaction:-
𝐴 + 𝐵 ↔ 𝐶+ 𝐷
we can express the rate law as: Rate=k[A]m[B]n
The exponents, ‘m’ and ‘n’ are the orders of the reaction and provide information on how the rate depends on the particular reactant. They can have integer values of 0, 1 or 2. (eg. If m is 0, then the rate does not depend on the concentration of A at all.
Determining rate laws allows us to predict the rate and the manner in which reactants disappear & products form. Additionally we can gain an insight into how the reaction may be occurring at the molecular level (i.e. the mechanistic steps involved). Reactions often occur through a series of individual steps which comprise the reaction mechanism.
NB. The values of the orders in the rate law MUST be determined by experiment and CAN NOT be determined from the stoichiometry of the reaction.
Beer’s Law & linearity
Those compounds that absorb UV light (eg those with conjugated double bonds) show an absorption
I = Io*10-(ecl)
Abs = log(Io/I) = ecl I0
profile in which a maximum wavelength (lmax) can be observed.
Beer’s Law states that the concentration of an analyte is related to its absorbance in a linear manner but only for dilute solutions. For more concentrated solutions the relationship is no longer linea
𝐴bs = 𝜀cI𝜀= molar absorptivity – a constant for a particular analyte
c = concentration
l = path length through which the UV beam travels through the analyte solution (= 1 cm cuvette)
Bleach and dyes:
Today you will be studying the kinetics of the reaction between the commercial bleach, sodium hypochlorite (NaOCl) and the food dye Fast Green FCF.
Fast Green FCF
MW 808.86
lmax 615 nm
Molar absorptivity e 1560
The reaction is:
NaOCl + Fast Green FCF à (Colourless products)
We can write a rate law expression for reactions 1 and 2:-
The objectives of the experiment are to determine the rate constant (k1), and the orders with respect to each reactant for each of the reactants (x, and y). Since this is a bimolecular reaction we use a ‘pseudo order’ expression for the rate law and ensure one reactant is in excess. Today this will be NaOCl since we can more easily measure colour changes. This way, its concentration is essentially constant.
We can define a new constant, kobs where kobs=k[NaOCl]x, and equation (3) becomes:-
𝑅ate=kobs [Green]y
All of the reactants and products are colourless except Fast Green FCF.
The [dye] can be determined by measuring the absorbance of the reaction mixture as a function of time.
Experimental Procedure
The procedure is quite simple – we are basically adding dye to the bleach and watching the color disappear. However, to determine the rate constants and orders we need to take quantitative measurements using a UV Spectrophotometer. We’ll also slow the reaction down to obtain reliable data.
You will find the following solutions at your bench:
Sodium hypochlorite solutions (20%, 10% & 5% dilutions of a 6% commercial solution)
10 mg/ml Fast Green
1 mg/ml Fast Green
0.1 mg/ml Fast Green
To begin we need to first decide on the best concentration of dye that will produce a linear relationship between absorbance & concentration. An absorbance range of 0.2-1.5 is generally accepted to be the linear range.
Draw up a table in your lab notebook that looks something like this:
Dye Used [Dye] 20% sodium
hypochlorite Reaction Time * (sec)
Fast Green 10 mg/ml
1 mg/ml
0.1 mg/ml
For this part, share the workload amongst other pairs of students on your bench.
Add 1 mL 20% sodium hypochlorite (bleach) to a 5mL tube
Add 100 µL of dye
Visually note the reaction time for the reaction to go to completion (clear)
*if the solution is still heavily coloured after ~10 minutes, then it is too concentrated.
Too fast a reaction rate will mean we can’t measure it accurately. (Too slow and we’ll be here all day!) A reaction time of ~10-15 minutes is ideal for making reliable measurements.
Select the concentration of dye that best satisfies Beer’s Law. (i.e. giving a reading between
~1-2 initially.)
Testing the rates
As mentioned to determine the rate constants & orders one reactant, the bleach, will be kept in excess with respect to the dye. (Nb. It’s easier to follow the extent of reaction by colour change so it must be the dye that is the limiting reagent).
In your lab notebook record the details in a table similar to that below. To determine the order for the reaction, we need to measure the reaction rate at different bleach concentrations. Start with a 20% dilution of the 6% commercial bleach.
A) Fast Green FCF
Time (s) Abst *[Dye]t ln[Dye]t 1/[Dyet]
0 0.895
30 0.808
60 0.728
90 0.651
120 0.579
150 0.516
180 0.459
210 0.409
240 0.364
270 0.323
300 0.287
330 0.256
360 0.227
390 0.203
420 0.181
450 0.162
480 0.144
510 0.130
540 0.117
570 0.105
600 0.094
(* Molar absorptivity for Fast Green is 1560)
*For these experiments distribute the workload amongst pairs of students on your bench.
Zero the spectrometer using a blank (1 mL 20% bleach diluted with 100µL distilled water). **If you change bleach concentrations you will need to repeat this step.
For the test solution, add 1 mL 20% bleach to a cuvette.
* You will be timing the reaction and taking absorbance every 30 seconds so make sure that the instant you add the dye be ready to quickly invert to mix, take the 1st measurement.
Add 100 µL Fast Green FCF to the cuvette. Invert quickly to mix & measure the absorbance at the appropriate l Start the timer.
Measure absorbance every 30 seconds for ~ 10 mins.
Repeat steps 1 – 4 with 10% bleach solution (dilution by half). Remember to re-zero your spectrometer with a new “blank” solution made from the new concentration of bleach.
Time (s) Abst *[Dye]t ln[Dye]t 1/[Dyet]
0 0.958
30 0.940
60 0.893
90 0.846
120 0.811
150 0.779
180 0.742
210 0.709
240 0.678
270 0.648
300 0.622
330 0.594
360 0.567
390 0.542
420 0.517
450 0.496
480 0.474
510 0.455
540 0.436
570 0.417
600
6.Repeat steps 1-4 with 5% bleach solution (make a 50% dilution of the 10%). Remember to re- zero your spectrometer with a new “blank” solution made from the new concentration of bleach.
Time (s) Abst *[Dye]t ln[Dye]t 1/[Dyet]
0 0.969
30 0.937
60 0.897
90 0.865
120 0.832
150 0.802
180 0.772
210 0.739
240 0.715
270 0.687
300 0.662
330 0.638
360 0.615
390 0.593
420 0.571
450 0.551
480 0.531
510 0.513
540 0.495
570 0.479
600 0.464
Results (Tables, graphs & calculations)
Your lab report will rely heavily upon the data that you have generated. You will be preparing a number of graphs and tables.
Remember that the slope of the straight line in your graphs of the integrated rate laws will determine the rate constant, kobs, and which of the integrated rate law graphs produces a straight line tells you the order with respect to the dye (y).
𝑅ate=kobs [Green]y
(where kobs = k[bleach]x since we’re keeping bleach in excess)
Reaction Order with respect to Dye: (by linearity of plots of integrated rate laws)
Plot 3 graphs to determine the order of the reaction, y, with respect to the dye:
[dye] vs time,
2. ln[dye] vs time,
3. 1/[dye] vs time.
Depending on which graph results in a straight line, you can tell whether the reaction is 0, 1st or 2nd order, respectively with respect to the dye, y.
Once you’ve determined which graph is linear, determine the slope of this graph. (= pseudo rate constant, kobs,) (Nb. for 0 or 1st order, the slope is equal to -kobs.)
Reaction Order with respect to Bleach: (by comparing effect on kobs when using difft [NaOCl]
Calculate kobs for the different [NaOCl] and compare these to find the order of the reaction with respect to (wrt) bleach, x:
2.
Expt. Reaction [NaOCl] (mM) kobs
1 1.61 (20% dilution)
2 0.81 (10% dilution)
3 0.40 (5% dilution)
If kobs remains constant for the different [NaOCl], the reaction is 0 order wrt NaOCl;
If kobs doubles as the [NaOCl] doubles, the reaction is 1st order wrt NaOCl
If kobs increases by factor of four as the [NaOCl] doubles, the reaction is 2nd wrt NaOCl.
When you’ve found the order with respect to bleach, x, you can now find the actual rate constant k, for the reaction, since
kobs = k [bleach]x
Use the values from one experiment (1, 2 or 3 above) to insert into the equation above to calculate the actual rate constant, k. (k = kobs/[bleach]x )
Now you can write the rate expression (rate law)
Once you have this, you can calculate the instantaneous rate for any combination of dye and bleach concentrations.
Discussion:
Refer to your aim – were they achieved?
Describe briefly your method for finding the optimum concentrations of dye and bleach with reference to Beer’s Law.
Discuss how the rate law you determined compares to what is known about the mechanism of the reaction of bleach and dye.
Talk about any deviation in your data: How close is your graph to linear? Do all three trials of each concentration of bleach and dye result in the same kobs? Does kobs increase/decrease by the same factor for each variation of concentration of bleach? What factors may have caused these deviations?
