Kinetics_v6c.docx
BleachReactionKinetics
The reaction between sodium hypochlorite and a blue dye is quantitatively monitored using optical
absorption spectroscopy.
1
1.1
OBJECTIVES
EXPERIMENTAL GOAL
Students will use absorbance spectroscopy to determine the rate law for the reaction of sodium
hypochlorite with a blue dye.
1.2
PREREQUISITE SKILLS AND KNOWLEDGE
Students are expected to have facility and experience with absorbance measurements on the SpectroVis
Plus Spectrophotometer.
1.3
RESEARCH SKILLS
After this lab, students will have had practice in:
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1.4
following laboratory protocols
using a laboratory notebook
following hazardous waste guidelines
using the proper personal protection equipment
choosing an appropriate size of micropipette
using micropipettes
calculating the molarity of diluted solutions
using a spectrophotometer to measure concentration
using a spectrophotometer to measure the course of a reaction in time
organizing data
using Excel to graphically analyze experimental data
LEARNING OBJECTIVES
After this lab, students will be able to:
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Use optical absorbance spectrometry to measure the rate of a reaction
Use experimental kinetics data to determine a rate constant
Use experimental kinetics data to determine the order of a reaction
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2|Reaction Kinetics
2
PRE-EXPERIMENT
In this experiment you will measure the timedependence of the reaction between sodium
hypochlorite, NaClO, and the blue dye Brilliant
Blue FCF (Blue #1). Blue #1 is usually in the form
of a disodium salt. The structure is shown in the
figure to the right.
The reaction involves two reactants:
NaClO + Blue #1  colorless products
You will measure the rate of disappearance of Blue
#1 to learn a little more about the reaction.
2.1
CHEMICAL REACTION RATES
A rate is a change in a variable over time. Reaction rate is usually measured as a change in reactant or
product concentration over time.
Thus, the rate of the general reaction A  B, shown graphically above, can be expressed as
Rate =
Δ[A]/Δt = Δ[B]/Δt
Here [A] is the concentration of the reactant A and [B] is the concentration of the product B. Why is there
a negative sign in front of Δ[A]/Δt?
As ∆ gets smaller, the average change in concentration over time approaches the time derivative,
d[A]/dt. The time derivative of the change in concentration is called the instantaneous reaction rate. The
instantaneous reaction rate is the slope of the curve shown above at a particular instant in time.
For the general reaction
aA + bB  cC + dD
(1)
the rate of reaction can be expressed as
1
1
1
Satisfy yourself that this is reasonable.
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2.1.1
The Rate Law
The rate law can be found experimentally, but cannot be predicted from the balanced equation or reaction
stoichiometry. The rate law describes how reactant concentration determines the rate of reaction. For the
general reaction (1), the rate law has the form
A
B
In this equation, k is the rate constant, which is specific to the reaction at a particular temperature. The
rate constant thus expresses how temperature affects the rate. The exponents m and n express the
“reaction orders”. They describe how the rate is affected by reactant concentration. If the rate doubles
when [A] doubles, then the reaction is said to be first order in A, and m = 1. If the reaction rate increases
by a factor of four when [B] doubles, then the reaction is said to be second order in B, and n = 2. If, on the
other hand, the rate does not depend on the concentration of B at all, then n = 0 for B.
2.2
REACTION MECHANISM
Brilliant Blue FCF is blue because it absorbs light in the visible spectrum. The highly conjugated* nature
of the chemical structure allows the molecule to absorb light at a wavelength that will promote an electron
into an excited state. This kind of absorption usually occurs at wavelengths in the visible or ultraviolet
range of the electromagnetic spectrum. A bonding structure is called conjugated when bonds alternate
between single and multiple bonds. You can see from the figure above that alternation between single and
double bonds occurs over most of the molecule. A disruption to any one of the double bonds in the
system can disrupt the conjugation and cause the molecule to stop absorbing visible light.
In aqueous solution sodium hypochlorite dissociates into sodium and hypochlorite ions:
NaClO  Na+ + OClSodium hypochlorite acts as a bleach because hypochlorite can react with a molecule containing double
bonds, through a number of different mechanisms, changing the double bond to a single bond and
disrupting the conjugated system. One possible mechanism is shown in the figure below. In this case,
OCl- facilitates the addition of water to the carbon near the center of the molecule. Look for the double
bond that changes to a single bond upon the addition of –OH to one carbon and –H to another.
*
If you want to know more, refer to http://en.wikipedia.org/wiki/Conjugated_system and/or a good organic
chemistry textbook.
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The conjugated pattern of alternating single and double bonds is broken at this carbon and the molecule
no longer absorbs light in the visible range.
Several possible mechanisms will give these results. Here are three for you to consider:
A.
OCl- + dye  OCldyeH2O + OCl-dye-  OCl- + H2O-dye
(slow)
(fast)
In the above mechanism, the hypochlorite ion forms a complex with the dye (OCldye-) that allows the
water to come in and break into the double bond. Or, perhaps two hypochlorite ions need to add to the dye
to open the double bond to the water molecule, as in mechanism B, below.
B.
dye + 2OCl-  OCldyeOCl
H2O + OCldyeOCl 2OCl- + H2O-dye
(slow)
(fast)
Compare A and B to mechanism C, in which water first forms a complex with the dye in the slow step.
The hypochlorite then moves in to break the double bond, which allows the water to be added across the
bond.
C.
2.2.1
H2O + dye H2Odye
H2Odye + OCl- OCl- + H2O-dye
(slow)
(fast)
Relating Reaction Order to Mechanism
Of greatest interest in a kinetics experiment is the slowest step. When you measure the rate of a reaction,
you are actually measuring the rate of the slowest step. This step is known as the rate-determining step or
rate-limiting step. In mechanism A the rate-limiting step is a reaction between one OCl- ion and one dye
molecule. It is first order in OCl- and first order in dye. The rate law for the rate-limiting step is thus
OCl
dye
Since the reaction cannot go any faster than the rate of the slowest step, the rate-limiting step determines
the rate of the reaction. If your measurements of reaction rate show that the reaction is first order in
bleach (OCl-) and also first order in dye, then your data support mechanism #1. If your measurements
show different orders, then a different mechanism is supported.
What orders of reaction would you need to measure for bleach and dye in order to support mechanism B?
What about mechanism C?
2.2.2
Integrated Rate Laws
Integrated rate laws allow you to describe the change in reaction rate over time.
Find the Excel file SampleKineticsData.xlsx in Student_Resources>Labs>05.Kinetics. Plot the sample
kinetics data according to the instructions below for each order of reaction. You should be able to
determine the order of the reaction for the absorbing reactant, and also the value of the rate constant k.
2.2.2.1 Zero-order reactions
For a reaction zero-order in A, the rate can be expressed as the derivative
and also in terms of the rate law
Setting these two expressions for rate equal gives the equation
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Integrating over time gives the integrated rate law for a zero-order reaction:
This equation can be rearranged into the form of an equation for a straight line in the y = mx + b form
For the above equation,
is y, t is x, –k is the slope (m), and
is the y-intercept (b).
To test whether the sample data is zero order in the absorbing reactant, you only need to plot the
Absorbance vs Time. If the data give a straight line, the reaction shows zero-order kinetics in that
reactant.
2.2.2.2 First-order reactions
For a reaction first-order in A, the rate is
And the integrated rate law is
ln
A0
At
ln
ln
The straight line for this equation is
ln
y
=
ln
mx
+
b
To test whether the sample data show first order kinetics, take the natural log (=LN(B2), etc.) of the
Absorbance points and plot these versus Time. If the resulting plot is a straight line the reaction shows
first order kinetics in the measured reactant.
2.2.2.3 Second-order reactions
For a reaction second-order in A, the integrated rate law is
1
At
1
A0
Rearranging gives the y = mx + b form:
1
A t
1
A0
To test whether the sample data show second order kinetics, take the inverse (=1/B2, etc.) of the
Absorbance points and plot these versus Time. If the resulting plot is a straight line, the reaction shows
second order kinetics in the absorbing reactant.
2.2.3
Simplifying a Complicated Reaction Order
Since the reaction you will be studying involves two reactants, the complete rate law for the reaction is
likely to include both reactants. In this case, the rate law will be
(2)
Blue #1
NaClO
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To simplify calculations, you will use sodium hypochlorite concentrations several times greater than the
concentrations of the dye. When the sodium hypochlorite is in excess, its concentration changes only by a
small negligible fraction. In this case, the product k[NaClO]n can be taken as a new constant
(3)
= k[NaClO]
and the observed kinetics describe only the change in concentration of the blue dye. The rate law in
equation (2) becomes
(4)
Blue #1
To learn the order of reaction with respect to sodium hypochlorite, measure the simplified rate constant
at several different concentrations of sodium hypochlorite. The order of the reaction with respect
versus [NaClO].
to sodium hypochlorite is the slope of a log-log plot of
Finally the rate constant k for equation (2) can be determined in a number of ways, such as by solving
and [NaClO].
equation (3) for k using several measured values of
2.3
OPTICAL ABSORBANCE MEASUREMENT
To follow the progress of this reaction, you will use the fact that the optical absorbance of a species is
linearly proportional to its concentration as long as the Beer-Lambert Law holds. The SpectroVis Plus
spectrophotometer, in addition to being able to display the entire absorption spectrum, can also be
configured to record the absorbance at a specific wavelength versus time.
As the bleaching reaction occurs, you will monitor the course of the reaction by recording the absorbance
at max for Blue 1 (approx. 630 nm). You will need to correct the measured absorbance (630 ) for
background by subtracting the absorbance at t =  (630
, where the absorbance has flatlined. You will
– 630
versus time in three different ways, as described above (2.2.2 Integrated Rate
then plot 630
Laws), in order to determine the order of reaction with respect to Blue 1.
2.4
PREPARE FOR THIS EXPERIMENT
Read through the entire laboratory manual as well as the resources provided under the Lessons tab.
Prepare a workbook in Excel for recording and analyzing your results. Email yourself a copy to use
during the lab.
Don’t forget to plot the sample kinetics data in the file SampleKineticsData.xlsx in
Student_Resources>Labs>05.Kinetics. This part of your assignment is described in more detail in section
2.2.2.
You will need to refer to both this and the previous absorbance laboratory manuals when preparing your
lab notebook for this experiment.
Before you begin this experiment, you should have a pretty good sense of what you will be doing. Make a
diagram or flow chart in your lab notebook to organize what you will be doing in the experiment.
When you feel ready, test your preparation using the Pre-Experiment Quiz.
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3
3.1
LABORATORY MANUAL
MATERIALS CHECK OFF LIST
Each small group of (2-3) students will have:
Laptop computer with LoggerPro software
Vernier SpectroVis Plus Spectrophotometer, or similar spectrophotometer
USB cable
4 disposable 4 mL cuvettes with tight-fitting lids (clean and reuse as needed)
Set of micropipettes
1000 L pipette tips
200 L pipette tips
10 or 25 mL volumetric flask
50 mL beaker
200 – 400 mL waste beaker
DI water bottle
Bench top waste bin for tips
Tube containing 5 mL 8.25% (w/v) NaClO
Tube containing 1 mL 3 x 10-4 M FD&C Blue #1
3.2
SAFETY AND WASTE DISPOSAL PROTOCOLS
You will be working with solutions of a strong oxidizer; goggles, closed-toe shoes, loose long pants or
long skirt, and a lab coat must be worn for this experiment.
Save used solutions in the waste beaker at your bench. Empty the waste beaker into the labeled waste
bottles when cleaning up.
3.3
3.3.1
EXPERIMENTAL PROCEDURE
Calibrate the Spectrometer
1. Connect the spectrophotometer to a USB port on the computer, open LoggerPro and check to make
sure the spectrophotometer is on.
Q1. How can you tell the spectrometer is on?
2. Fill your sample cuvette with 3000 µL of DI water and insert it in the sample holder. This is your
reference cuvette.
3. When the spectrophotometer has warmed up, calibrate it.
3.3.2
Measure Reference Spectra of Reactant and Product
1. Add 50 L of the stock (3.0 x 10-4 M) Blue 1 solution to the cuvette. Place a clean lid over the top of
the cuvette, and press it firmly in place with the pad of one finger. Make sure the top is sealed. Using
your finger to hold down the lid, mix the dye and water by inverting the cuvette two or three times. If
any solution leaked out the side, wipe the cuvette clean with a lab wipe.
NOTE: If you are unsure of how to safely accomplish this task, ask your instructor for help.
Make sure you can do this quickly and comfortably without spilling; you will be doing the same thing
with bleach in the cuvette.
2. Record the absorption spectrum of the unreacted Blue 1. Save this spectrum for future reference.
Make sure your peak absorbance is in the linear range for Blue 1.
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3. Click on the spectrum at the peak maximum so that you can record the absorbance and max in your
notebook.
4. Pipette 100 L of 8.25% NaClO into the cuvette cell, cover with a lid, and mix by inverting as
described above. Return the cuvette to the sample holder. Try to complete this step within 30 seconds.
5. Observe the reaction in real time in absorbance mode. Note how the reactant peak diminishes and
eventually disappears into the baseline.
Q2. What is the concentration in M of Blue #1 in the cuvette the instant the bleach is added? How did
you calculate this value?
Q3. The molar mass of sodium hypochlorite (NaClO) is 74.44 g. What is the concentration in M of
NaClO in the cuvette the instant after the bleach is added? Confirm your calculation with your
extended group.1
3.3.3
Kinetics Measurements: Absorbance over Time
For the kinetics measurements, you will repeat the above procedures, starting with the reference
spectrum of the cuvette filled with water, but you will record the change in absorbance over time.
Set up the spectrophotometer to record Absorbance vs. Time.
3.3.3.1 Generate a Spectrum
1.
2.
3.
4.
5.
Place a fresh sample of 3000 µL DI water into the cuvette.
Calibrate the spectrometer.
Add 50 µL Blue #1 solution to the cuvette and mix.
Place the cuvette into the spectrometer and click >Collect.
Click Stop to end data collection
3.3.3.2 Configure SpectroVis Plus to Measure Absorbance vs. Time
1. Click on the Configure Spectrophotometer Data Collection
button.
2. Select Absorbance vs. Time as the data-collection mode. The default wavelength is max. Click OK
to continue at this wavelength, or Clear to select a different wavelength.
Adjust the sampling parameters by choosing Data Collection from the Experiment menu.
3. Choose a duration and sampling rate so that the spectrometer will record an absorbance once every
2.5-5 seconds for 5 minutes. Record in you lab notebook the sampling rate and duration.
Q4. How many points will you collect for each 5 minute run? Show your calculations.
4. Mix the reactants as before, place the cuvette in the spectrometer and click >Collect.
5. Record for about 5 minutes, or until the absorbance flatlines. Click Stop to stop data collection if the
absorbance flatlines before the 5 minutes are up.
To measure the rate of reaction at different bleach concentrations, repeat the kinetics measurement using
a range of different aliquots of NaClO. Choose NaClO amounts over as large a range of values as
possible that still give reaction rates in a conveniently observable timescale. Adjust the data acquisition
parameters appropriately to produce the most complete and accurate data set.
Note that most runs will reach the flatline stage within about 2 minutes.
3.3.4
Data Analysis
3.3.4.1 Order of reaction with respect to Blue #1
Use the integrated rate laws to determine the order of reaction for the Blue #1.
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Hints for successful data analysis:
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Do not forget to subtract the background absorbance from your data: A = 630
Do not use data from the flatline region. Why not?
– 630
Q5. Paste a snip of the successful plot here.
Q6. What did you get for the order of the reaction with respect to the concentration of Blue #1?
3.3.4.2 Order of reaction with respect to sodium hypochlorite
Recall from section 2.2.3 that when the concentration of sodium hypochlorite NaClO is in excess, its
concentration can be considered as constant over the length of the bleaching reaction.
Q7. What is the simplified rate law when you use this assumption?2
Q8. What is the simplified rate constant?3
Using the same integrated rate law that you used to find the order of reaction for the blue dye, determine
for each concentration of sodium hypochlorite you tested. Use this
the simplified rate constant
information to find the order of reaction with respect to sodium hypochlorite.
Q9. What is the order of reaction with respect to sodium hypochlorite?
Q10. In a few complete sentences, describe the method you used to achieve the answer to the previous
question. Use graphs as necessary.
3.3.4.3 The rate law for the reaction of Blue #1 and sodium hypochlorite
Determine the value of the constant k for the rate law that describes the reaction of Blue 1 with sodium
hypochlorite.
Q11. In a few complete sentences describe the method you used to find the rate constant for this reaction.
Use graphs as necessary.
Q12. Write the rate law for the reaction here.
3.4
POST-LAB ASSIGNMENT
Work with your extended group to submit a one-page abstract in class, describing the experiment you just
completed.
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