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CHEMISTRY HIGHER LEVEL INTERNAL ASSESSMENT
AN INVESTIGATION TO DETERMINE THE RATE LAW OF THE
CHEMILUMINESCENT OXIDATION OF LUMINOL BY SODIUM HYDROXIDE AND
HYDROGEN PEROXIDE
RESEARCH QUESTION What is the rate law of the chemiluminescent
oxidation of luminol (C8H7N3O2 ) by sodium hydroxide (NaOH ) and
hydrogen peroxide (H2O2)? AIM OF THE INTERNAL ASSESSMENT The aim of
this experiment is to experimentally derive the rate law for the
chemiluminescent oxidation of luminol by determining the reaction
orders with respect to each reactant. This will be done by varying
the concentrations of each reactant and determining the effect this
has on the rate of photon emission, as measured by the change in
illuminance of the reaction mixture over time. INTRODUCTION I found
myself initially interested in exploring chemical kinetics after a
conversation with a classmate about the real-life applications of
the field. After some discussion, we came to appreciate the large
scope of applications of chemical kinetics, particularly in
phenomena which we would encounter on a day to day basis such as
the corrosion of metal and the shelf life of food products.
Following this conversation, I began further researching chemical
kinetics when I encountered an application which particularly
excited me; the chemiluminescence of luminol. As somewhat of a
crime drama fanatic, I was very familiar with the spray used by
forensic investigators to make blood almost magically glow in crime
scenes, yet I never really understood the science behind this
process. The chemiluminescence of luminol was exactly this; a
reaction catalyzed by the heme iron in blood which would emit a
distinct glow and thus allow the blood tracings in crime scenes to
be 1visualized. Ultimately, I chose to use this investigation to
explore the kinetics behind this reaction, specifically
scrutinizing the effect which the concentration of each reactant
has on the reaction rate while the ultimate goal of determining the
order of reaction with respect to each reactant and, subsequently,
the rate law for the reaction. BACKGROUND INFORMATION Within
chemistry there exists the field of chemical kinetics which is
concerned with “the effects of different factors on the rates of
reaction, reaction mechanisms and forming models to predict
reaction rates” . One of 2the elementary concepts of chemical
kinetics is the “rate” of a reaction, defined as the “change in the
concentration of the reactants or products of a reaction per unit
time” . To explain why different reactions occur at different
3rates the collision theory is utilized, which is composed of three
principles that must be true for a reaction to occur . The first
principle of the collision 4theory is that reactant particles must
collide. The second principle of the collision theory is that
reactant particles must collide with the correct orientation in
order to “allow atoms that bond together to be in contact with one
another” . The last principle of the collision theory is that
particles must 5collide with sufficient energy that is equal or
greater than the activation energy - “the minimum energy which
particles must possess in order to break pre-existing bonds in the
reactants and form new bonds in the products” . Critically, any
factor which influences the rate of a reaction must 6influence one
of these defining principles . 7
1 “NCATS Inxight: Drugs - LUMINOL.” Inxight Drugs. US Department
of Health & Human Services. Accessed November 24, 2019.
https://drugs.ncats.io/drug/5EXP385Q4F.
2 Helmenstine, Anne Marie. “Understand Chemical Kinetics and
Rate of Reaction.” ThoughtCo. ThoughtCo, March 2, 2019.
https://www.thoughtco.com/definition-of-chemical-kinetics-604907.
3 “Reaction Rate.” Chemistry LibreTexts. Libretexts, September
30, 2019.
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Kinetics/Reaction_Rates/Reaction_Rate.
4 “The Collision Theory.” Chemistry LibreTexts. Libretexts,
September 30, 2019.
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Kinetics/Modeling_Reaction_Kinetics/Collision_Theory/The_Collision_Theory.
5 “Collision Theory.” Chemistry. PressCorp. Accessed November
24, 2019.
https://opentextbc.ca/chemistry/chapter/12-5-collision-theory/. 6
“Activation Energy.” Definition of Activation Energy | Chegg.com.
CheggStudy. Accessed November 24, 2019.
https://www.chegg.com/homework-help/definitions/activation-energy-6.
7 “The Collision Theory.” Chemistry LibreTexts. Libretexts,
September 30, 2019.
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Kinetics/Modeling_Reaction_Kinetics/Collision_Theory/The_Collision_Theory.
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One example of a factor which influences the rate of a reaction
is the concentration of the reactants. Increasing the
concentration of the reactants increases the molar amount of
reactant particles per unit volume, thus increasing the frequency
of successful collisions between particles and hence the rate of
the reaction. Moreover, an increased concentration will result in a
greater number of reactant particles possessing energy greater or
equal to the activation energy, further causing the frequency of
successful collisions to increase and thus increasing the rate of
the reaction. This phenomenon is illustrated in the Boltzmann
distribution curve in Figure 1. As can be seen, while an increased
concentration of the reactants (c1) doesn’t alter the shape of the
Boltzmann curve, meaning that the mean kinetic energy of particles
remains the same, the area under the c1 curve is greater, resulting
in a greater number of particles possessing energy that is larger
or equal to the activation energy (Ea). The relationship between
the rate of a reaction and the concentration of its reactants can
be expressed in a rate law , for example: 8
For the hypothetical reaction , the rate law can be expressed as
A B Ca + b → c ate k[A] [B]r = x y
In the above rate law, [A] and [B] represent the concentrations
of the two reactions in mol dm-3. The constant k is known as the
rate constant and is a coefficient of proportionality which
represents the relationship between the rate of the reaction and
the concentration of the reactants . The rate constant is specific
for the reaction at a particular temperature and its units 9
10depend on the number of reactants in the rate equation and the
units expressing the rate of the reaction. The two coefficients, x
and y, represent the order of the reaction with respect to each
reactant. These coefficients don’t necessarily correspond to the
stoichiometric coefficient in the reaction equation . The order of
reaction with respect to a reactant is the “power dependency” 11of
the rate on the concentration of the reactant . Commonly, the order
of reaction with respect to a reactant will be zero-order,
12first-order or second-order. The zero-order reactant is raised to
the zeroth power, meaning that the reaction rate is independent of
that reactant’s concentration. A first-order reaction is raised to
the first power, meaning that the reaction rate is directly
proportional to its concentration. A second-order reactant is
raised to the second power, meaning that the relationship between
the reaction rate and the reactant’s concentration is exponential.
If these different orders of reaction were to be illustrated on a
concentration against rate graph, a zero-order reaction will be a
horizontal line, a first-order reaction will be a linear graph
starting at the origin and a second-order reaction will be an
exponential graph starting at the origin.
The reaction explored in this investigation is the
chemiluminescent oxidation of luminol. This reaction is exothermic,
meaning that the energy of the reactants is greater than the energy
of the products and there’s thus a net release of energy by the
reaction . However, unlike most exothermic reactions that release
energy in the form of heat, this reaction releases 13energy in the
form of photons, or light. This is what the term
“chemiluminescence” refers to; the emission of cold light (light
accompanied by little or no heat ) due to the relaxation of an
excited state . The reactant which is primarily responsible for 14
15the chemiluminescent characteristics of this reaction is luminol
(formally known as 5-amino-2,3-dihydro-1,4-phthalazinedione ), an
organic compound that exhibits chemiluminescence when activated by
an 16oxidizing agent in an alkaline solution in the presence of a
catalyst . In this specific reaction, the oxidizing agent is
hydrogen 17peroxide (H2O2), the base is sodium hydroxide (NaOH) and
the catalyst is potassium ferricyanide (K3[Fe(CN)6]). The full
reaction is represented by the equation below.
H N O O NaOH H NO H O Na v C8 7 3 2 (s) + 2H2 2 (aq) + 2 (aq) →
C8 5 4
2−(aq) + N 2 (g) + 4 2 (l) + 2
+(aq) + h
One of the products of this reaction, hv, represents the energy
of the light emitted by the reaction. The particle-wave
duality of light allows for the energy of a light ‘particle’ (a
photon) to be expressed as a function of the light wave’s
frequency. This relationship is expressed by Planck’s equation: ,
which states that the energy of a photon is equal E v"" = h 18
to the product of the frequency of the light wave (v) and
Planck’s constant (h).
8 “The Rate Law.” Chemistry LibreTexts. Libretexts, June 5,
2019.
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Kinetics/Rate_Laws/The_Rate_Law.
9 Ball, David W., and Jessie A. Key. “Rate Laws.” Introductory
Chemistry 1st Canadian Edition. BCcampus, September 16, 2014.
https://opentextbc.ca/introductorychemistry/chapter/rate-laws-2/.
10 Ibid. 11 Ibid. 12 Bagshaw, Clive R. 2013. “Order of
Reaction.” In Encyclopedia of Biophysics, edited by Gordon C. K.
Roberts, 1807–8. Berlin, Heidelberg: Springer
Berlin Heidelberg.
https://doi.org/10.1007/978-3-642-16712-6_575. 13 Firuz, Akmal.
"Chemiluminescence." n.d. PPT. 14 “Cold Light.” The Free
Dictionary. Farlex. Accessed November 24, 2019.
https://www.thefreedictionary.com/cold light. 15 “Luminol.”
Chemistry LibreTexts. Libretexts, June 5, 2019.
https://chem.libretexts.org/Bookshelves/Analytical_Chemistry/Supplemental_Modules_(Analytical_Chemistry)/Analytical_Chemiluminescence/2:_Chemiluminescence_Reagents/2.01:_Luminol.
16 Clegg, Brian. “Luminol.” Chemistry World, April 16, 2014.
https://www.chemistryworld.com/podcasts/luminol/7272.article. 17
Lodovico, Ray. “Luminol and Chemiluminescence.” PhysicsOpenLab,
February 6, 2019. http://physicsopenlab.org/2019/02/06/luminol-2/.
18 “E=Hv.” PhysicsMatters.org: Quantum: E=hv. American Physical
Society, 2012. http://www.physicsmatters.org/quantum/ehv.html.
2
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In order to get a better understanding of the role which each
reactant plays in the reaction, and how the reaction is able to
emit light, the reaction mechanism must be explained. The
proposed mechanism for the reaction is shown in Figure 2. It is
important to note that the initial presence of H2O2, NaOH and the
catalyst (K3[Fe(CN)6]) is omitted in this mechanism. In the
reaction, each luminol molecule reacts with two OH- ions to form a
dianion, which is labelled (A) in Figure 2. The OH- ions are
produced by the dissociation of 1 mole of NaOH in the solution into
1 mole of Na+ ions and 1 mole of OH- ions, as illustrated by the
following equation : 19
aOH aN (aq) → N +(aq) + OH
− (aq)
To form the dianion (A), each OH- ion donates a pair of
electrons and reacts with a hydrogen atom in luminol to form a
water molecule. This is accompanied by the heterolytic fission of
the N-H bond in luminol, resulting in each nitrogen atom acquiring
a negative charge (hence the name dianion). Dianion (A) is a
resonance structure, as is evident by its adjacent single and
double bonds, and is thus stabilized by its delocalized electrons .
The other resonance structure of dianion (A) is 20labelled (B) in
Figure 2. The resonance dianion (B) subsequently reacts with an O2
molecule produced by the decomposition of hydrogen peroxide, as is
illustrated in the equation below : 21
O H O2H2 2 (aq) → 2 2 (l) + O2 (g)
The reaction between dianion (B) and the O2 molecule forms the
cyclic peroxide (C), which directly converts into the
3-aminophthalate* transition state by releasing two nitrogen atoms
as nitrogen gas (N2) . The asterisk on the
22
3-aminophthalate* molecule indicates that is in an excited
state, meaning that the electrons in the molecule have become
promoted from the ground state to a higher energy or “excited”
state. This excited state is, however, very unstable, and the
electrons in the 3-aminophthalate* quickly fall back down from the
excited state to a lower energy state, emitting a quantized amount
of energy in the form of photons in the process . This causes the
transition state to decay and form the product (D). 23Ultimately,
the light emitted by the decayed transition state has energy hv and
has been shown to have a “wavelength of 425 nm, which appears in
the blue zone of the visible spectrum” . 24
Apart from the interesting theory behind this reaction,
investigating the rate law for the oxidation of luminol also has
real-life significance. One of the best-known uses of luminol is in
the forensic identification of blood traces at crime scenes . 25A
luminol solution containing a base (NaOH) and an oxidizing agent
(H2O2) is sprayed on a surface to be tested. Instead of potassium
ferricyanide, the heme iron in blood catalyzes the oxidation
reaction, meaning that any traces of blood will be instantly
detectable by the characteristic chemiluminescence of the luminol
solution . Ultimately, by investigating the 26kinetics of the
chemiluminescent oxidation of luminol, it will be possible to
suggest the most efficient and effective way to maximize the rate
of luminol’s chemiluminescence . For example, if it is found that
this reaction is second order with 27respect to one of the
reactants, it would be most effective to increase the concentration
of that reactant in order to increase the rate of luminol’s
oxidation.
19 “Definitions of Acids and Bases, and the Role of Water.”
Bodner Research Web. Accessed November 24, 2019.
https://chemed.chem.purdue.edu/genchem/topicreview/bp/ch11/acidbase.php.
20 Fleming, Declan. “Mechanism of the Reaction.” The
Chemiluminescence of Luminol - Home. Accessed November 24, 2019.
http://www.chm.bris.ac.uk/webprojects2002/fleming/mechanism.htm.
21 “The Catalytic Decomposition of Hydrogen Peroxide, II.”
Chemed. Accessed November 24, 2019.
https://chemed.chem.purdue.edu/demos/main_pages/19.7.html.
22 Fleming, Declan. “Mechanism of the Reaction.” The
Chemiluminescence of Luminol - Home. Accessed November 24, 2019.
http://www.chm.bris.ac.uk/webprojects2002/fleming/mechanism.htm.
23 “Forensic Analysis of Blood Using Luminol.” Faculty Virginia.
Accessed November 24, 2019.
http://faculty.virginia.edu/analyticalchemistry/Luminol/LUMINOL.html.
24 “Chemiluminescence - A Toast to Chemistry.” Flinn Scientific
Canada. Accessed November 24, 2019.
https://www.flinnsci.ca/chemiluminescence--a-toast-to-chemistry/Document/?contentId=7db395b8-885a-452e-b32d-5b7397054119.
25 Firuz, Akmal. "Chemiluminescence." n.d. PPT. 26 “NCATS
Inxight: Drugs - LUMINOL.” Inxight Drugs. US Department of Health
& Human Services. Accessed November 24, 2019.
https://drugs.ncats.io/drug/5EXP385Q4F. 27 Laidler, Keith J.
“Chemical Kinetics.” Encyclopædia Britannica. Encyclopædia
Britannica, inc., May 16, 2013.
https://www.britannica.com/science/chemical-kinetics. 3
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HYPOTHESIS From the reaction mechanism shown in Figure 2 it can
be predicted that both luminol and NaOH are part of the rate law,
given that they are both required to form dianion (A) which
proceeds to react with O2. However, if the reaction between dianion
(A) and the O2 molecule is the slow, rate-determining step then O2,
and thus H2O2, may also be part of the rate law. Given this, it is
necessary to use a method to predict the rate law. The steady-state
approximation is used to estimate the rate laws of reactions by
assuming that “the concentration of all reaction intermediates
remains constant throughout the experiment” . Another assumption
made as part of the steady-state approximation is that the “overall
rate law will be the 28final step in the reaction mechanism” . In
order to better understand the reaction, it is easiest to
deconstruct the reaction 29mechanism presented in Figure 2 into
elementary steps. For simplicity’s sake, I have drawn out some of
the compounds which take part in the reaction, as indicated in the
reaction mechanism, and have labelled them accordingly in Figures 3
to 6.
Figure 3: Luminol (I) Figure 4: Intermediate 1 (II) Figure 5:
Intermediate 2 (III) Figure 6: Product (IV)
As such, the following reaction mechanism was constructed,
following the findings of White et al. : 30
Given the steady-state approximation, the concentration of
intermediate 1 (II) and intermediate 2 (III) will remain constant
throughout the reaction, meaning that they will be produced at the
same rate that they are consumed by the reaction. Using the
reaction mechanism outlined above, it is possible to express the
rate of production and consumption of each intermediate in the form
of a rate law. These rate laws will be equal to each other and, via
some basic arithmetical operations, an expression can be deduced
for the concentration of each intermediate in terms of the rate
constants of other elementary steps and the concentration of the
original reactants. For example, the expression for the
concentration of intermediate 1 (II) is derived in the following
way:
ate of production of Ⅱ rate of consumption of Ⅱr = [Ⅰ][OH ]
[Ⅱ][OH ]k1
− = k2−
Ⅱ][ = k [OH ]2 −k [Ⅰ][OH ]1
−
= k2k [Ⅰ]1
Similarly, an expression for the concentration of intermediate 2
(III) can be derived. Given that III is involved in the equilibrium
step 2, its rate of consumption will be equal to the sum of the
rate of the backwards reaction in step 2 and the rate of step
3.
ate of production of Ⅲ rate of consumption of Ⅲr =
[Ⅱ][OH ] [Ⅲ][H O] [Ⅲ][O ]k2− = k−2 2 + k3 2
[Ⅱ][OH ] Ⅲ](k [H O] [O ])k2− = [ −2 2 + k3 2
28 Mott, Vallerie. “Steady-State Approximation.” Introduction to
Chemistry. Lumen. Accessed November 24, 2019.
https://courses.lumenlearning.com/introchem/chapter/steady-state-approximation/.
29 Ibid. 30 White, Emil H., Oliver. Zafiriou, Heinz H. Kagi, and
John H. M. Hill. 1964. “Chemiluminescence of Luminol: The Chemical
Reaction.” Journal of the
American Chemical Society 86 (5): 940–41.
https://doi.org/10.1021/ja01059a050. 4
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Ⅲ][ = k [Ⅱ][OH ]2
−
(k [H O]+k [O ])−2 2 3 2
Ultimately, the overall rate law can be generated by
substituting the derived values for the concentrations of
intermediate 1 (II) and intermediate 2 (III). The assumption that
the overall rate law is equal to the rate law of the final step
(step 3) is illustrated here.
When generating this overall rate law, two further assumptions
were made. The first assumption is that the value for k-2 is much
larger than the value of k3 ( ), meaning that the value of k3[O2 ]
is negligible in comparison to the value of >k−2 > k3 k-2[O2
]. This assumption based on the fact that “there is a lot more
water in an aqueous solution than there is O2 from the
decomposition of H2O2” . Ultimately, this assumption allows us to
remove k3[O2 ] from the rate law. The second assumption
31
is that water has “unit activity”, meaning that its
concentration has no effect on the reaction rate . This is because
water acts 32as a solvent in this reaction and its concentration is
therefore “extremely large and virtually constant” . This allows us
to 33
remove the concentration of water from the rate law. Ultimately,
a new rate constant ( k’ ), is formed from , making the k−2k k3
1
overall rate law equal to rate = k’[I][OH-][O2]. Critically, the
steady-state approximation suggests that the chemiluminescent
oxidation of luminol is first order with respect to luminol, OH-
ions and O2. This will act as my hypothesis for this investigation.
APPROACH TO THE RESEARCH QUESTION When considering the research
question, it was clear to me that I needed to construct a method
which would alter the concentrations of each reactant and measure
the rate of the luminol reaction at different combinations of
reactant concentrations. This would ultimately allow me to
construct a graph of concentration against reaction rate for each
reactant to determine its order of reaction. Initially, I needed to
determine a suitable way to measure the rate of the reaction. Given
the characteristic chemiluminescence of the reaction in question,
it was evident to me that calculating the reaction rate as the
change in illuminance of the reaction mixture over the change in
time would be most suitable. Illuminance, measured in lux (lx), is
a measure of total luminous flux per unit area and is proportional
to light intensity (measured in candela) . 34In this reaction,
concentration can be considered proportional to illuminance given
that, as concentration increases, the rate of the reaction
increases, leading to the formation of more unstable
3-aminophthalate* transition states per unit time. The greater the
number of unstable transition states, the more will decay into a
more stable state and emit light per unit time. This will translate
into a greater intensity of light emitted by the reaction mixture
and hence, given that illuminance is proportional to light
intensity, a greater illuminance value.
I wasn’t able to find any online methodology that determined the
rate law for this specific reaction so I decided to adapt
pre-existing methodologies that dealt with the oxidation of luminol
such that I was able to alter the concentrations of the reactants
and determine the rate law. I ultimately chose to alter the
methodology proposed by the department of chemistry at the North
Carolina State University (NCSU) , which instructed the use of two
solutions 35which would be mixed together to initiate the reaction.
In preliminary testing, I adapted this methodology by conducting
the experiment inside a box to block external light and by using a
light meter to measure the change in illumination of the reaction
mixture over time. By graphing these two variables against each
other, I was able to identify the spike in illuminance caused by
the emission of light once the reaction was initiated. By
subtracting
31 Fleming, Declan. “Mechanism of the Reaction.” The
Chemiluminescence of Luminol - Home. Accessed November 24, 2019.
http://www.chm.bris.ac.uk/webprojects2002/fleming/mechanism.htm.
32 “Rate Laws: An Introduction.” Chemistry LibreTexts.
Libretexts, April 27, 2019.
https://chem.libretexts.org/Bookshelves/General_Chemistry/Map:_Chemistry_(Zumdahl_and_Decoste)/15:_Chemical_Kinetics/15.2_Rate_Laws:_An_Introduction.
33 “Acid and Base Ionization Constants.” Chemistry for
Non-Majors. Lumen. Accessed November 24, 2019.
https://courses.lumenlearning.com/cheminter/chapter/acid-and-base-ionization-constants/.
34 Keim, Robert. “Understanding Illuminance: What's in a Lux? -
Technical Articles.” All About Circuits, January 11, 2016.
https://www.allaboutcircuits.com/technical-articles/understanding-illuminance-whats-in-a-lux/.
35 "Chemiluminescence with Luminol." North Carolina State
University Department of Chemistry. n.d. Raleigh. PDF. 5
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the illuminance prior to the spike (the initial illuminance)
from the maximum illuminance at the peak of the spike (the final
illuminance) I would be able to determine the change in
illuminance. Additionally, by subtracting the time at which the
initial illuminance was recorded (initial time) from the time at
which the final illuminance was recorded (final time) I would also
be able to determine the change in time. Ultimately, I could divide
the change in illuminance by the change in time in order to
calculate the rate of the reaction in illuminance over time (lx
s-1). A section of a graph picturing a spike in illuminance and the
points of the initial and final time and illuminance is shown in
Figure 7.
The methodology proposed by the NCSU also suggested the
appropriate concentrations and amounts of each reactant in these
solutions. I tested out solutions with differing concentrations of
each reactant in preliminary testing and immediately faced an issue
regarding the methodology. No matter the combination of reactants
in the two solutions, there would always be the risk of an
unintended reaction occurring. For example, if luminol and NaOH
were in the same solution, they would react preliminarily to form
luminol dianions. Moreover, if NaOH and H2O2 were in the same
solution, they would react preliminarily to form sodium peroxide .
Ultimately, I chose to mix NaOH and luminol in the same solution in
order to 36preserve the order of elementary steps in the reaction
mechanism (shown in Figure 2). The other reactant solution would
therefore contain H2O2 and K3[Fe(CN)6]. Evidently, the limitation
of this methodological approach is that the effects of the
concentration of OH- ions couldn’t be tested, as all the OH- ions
would preliminarily react to form luminol dianions. Hence, the rate
law which would actually be investigated in this experiment would
be:
ate luminol dianions] [O ]r = [ x 2
y Given that the molar ratio of luminol to luminol dianions in
the reaction with sodium hydroxide is 1:1 (as is illustrated in
Figure 2), the concentration of luminol dianions in the solution
should be equal to the concentration of luminol. Hence, when
applying the predicted first order with respect to both luminol and
oxygen from the steady-state approximation, the hypothesized rate
law becomes:
ate luminol][O ]r = [ 2 VARIABLES The independent variable in
this experiment is the concentration of the reactant solutions
used. As explained previously, two different solution types are
used in this experiment. These will be called the alkaline luminol
solution (contains distilled water, luminol and sodium hydroxide)
and the oxidizing solution (contains distilled water, hydrogen
peroxide and potassium ferricyanide). By varying the amounts of
luminol, sodium hydroxide and hydrogen peroxide in these solutions,
their concentrations can be altered and tested. In this
investigation, the solutions will be made such that a range of
three different concentrations is achieved for each of the
reactants in order to graph an observable trend and determine the
order of reaction with respect to each reactant. The dependent
variable of this experiment is the rate of the reaction, as
calculated by the change in illuminance over time (lx s-1). To
enhance the precision of the rate values calculated, 3 trials will
be conducted for each run of the experiment. The recording of the
time and illuminance will begin 5 seconds prior to the two
solutions being mixed in order to record a baseline illuminance
value before the reaction is initiated and to prevent any errors
caused by a potential time delay.
Controlled variables: 1. Temperature: Temperature is another
factor which influences the rate of a reaction. At higher
temperatures, reactant
particles have a higher average kinetic energy. This doesn’t
only mean that a greater number of particles possess energy that is
greater or equal to the activation energy, but this increased
temperature also causes particles to move faster and thus collide
more frequently , both of which increase the rate of the reaction.
Consequently, it is necessary to control 37the temperature of the
surroundings, as varying temperatures will increase/decrease the
rate of the reaction. To ensure this variable remains constant, the
experiment was conducted in the same laboratory room at a
temperature of 22.3 ºC. This temperature was checked prior to and
controlled throughout the experiment using a thermometer.
2. Light conditions: Given that illuminance is measured in this
investigation, it is necessary to conduct the experiment in a dark
space to ensure that the only change in illuminance recorded is due
to the reaction and not from external light sources. To ensure this
variable remains constant, it was chosen to conduct the experiment
in a light-blocking box (as was mentioned in the APPROACH TO THE
RESEARCH QUESTION section). This light-blocking box will be
constructed with two openings, a light sensor opening for the light
sensor and a beaker opening for the beaker (the reaction vessel) to
be placed into and taken out of the box. The box will also be
covered with a light-blocking material
36 "FACE Report No. 09CA004, A Chemist Dies from Burns Caused by
Mixing Chemicals." California Department of Public Health, May 28,
2009. Richmond. PDF.
37 “Factors That Affect Reaction Rates.” Chemistry LibreTexts.
Libretexts, June 5, 2019.
https://chem.libretexts.org/Bookshelves/Physical_and_Theoretical_Chemistry_Textbook_Maps/Supplemental_Modules_(Physical_and_Theoretical_Chemistry)/Kinetics/Reaction_Rates/Factors_That_Affect_Reaction_Rates.
6
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(black, opaque bin bag material) in order to prevent external
light from passing into the box through any pre-existing holes or
cavities.
3. Distance of the reaction vessel from the light sensor: The
intensity of light at different distances from a light source is
described by the inverse-square law . Critically, the intensity of
light is inversely proportional to the square of the 38distance,
meaning that the light intensity decreases as the distance from the
light source increases. Given that illuminance is proportional to
light intensity, the distance of the reaction vessel (the “light
source”) from the light sensor will affect the illuminance value
recorded and must therefore be controlled. To ensure this variable
remains constant, a circular marking will be made on the inside of
the light-blocking box to indicate where the reaction vessel (a 50
cm3 beaker) should be placed during each experimental trial. This
circular marking will be made to ensure that the closest face of
this beaker will be an insignificant 1 cm away from the light
sensor. This was measured using a 30 cm ruler.
4. Mass of potassium ferricyanide catalyst used: In this
reaction, potassium ferricyanide acts as a catalyst to speed up the
rate of the reaction. In its processed form, potassium ferricyanide
exists as small salt granules. Ultimately, a larger mass of
potassium ferricyanide would increase the total surface area of the
catalyst used, meaning that collisions between the catalyst and the
reactant particles would be more frequent and hence the rate of
reaction would increase. Critically, this illustrates why the mass
of potassium ferricyanide used needs to be controlled. To ensure
this variable remains constant, the same mass of potassium
ferricyanide (3.00 grams) was measured using an electronic weighing
scale and added to each oxidizing solution, as recommended by the
NCSU’s methodology . 39
5. Volume of alkaline luminol solution and oxidizing solution
mixed: Larger volumes of both the alkaline luminol and oxidizing
solutions would contain a greater number of moles of the reactants
used. Given this, larger volumes of, for example, the oxidizing
solution would result in more light being produced by the reaction
and hence a higher illuminance value being recorded. To ensure this
variable remains constant, the volume of the alkaline luminol and
oxidizing solutions was kept constant at 10 cm3 (meaning that 10
cm3 of each solution were mixed together to initiate the reaction).
This was measured using a 10 cm3 volumetric pipette (used for the
alkaline luminol solution) and a 10 cm3 syringe (used for the
oxidizing solution).
APPARATUS WITH UNCERTAINTIES
(1) light sensor ± 0.0001 lx (1) 10 cm3 syringe ± 0.5 cm3
(6) 250 cm3 volumetric flask ± 0.12 cm3 (1) 10 cm3 volumetric
pipette ± 0.03 cm3
(1) electronic weighing scale ± 0.005g (1) 10 cm3 graduated
pipette ± 0.05 cm3
METHODOLOGY The method for this experiment can be separated into
five parts: A) PART 1: constructing the apparatus 1. Take a
cardboard box and cut two openings into it using a pair of
scissors, as illustrated in Diagram 1. One of the
openings, the beaker opening, should be a 6 x 6 cm square on the
top of the box that is 12 cm from either edge of the box. The other
opening, the light sensor opening, should be a circle that’s 2 cm
in diameter whose centre is 1 cm from the bottom edge of the box
and 15 cm from the side edges of the box.
2. It is necessary to make a circular marking inside the box to
act as an indication of where to place the 50 cm3 beaker (with a
diameter of 5 cm) during each experimental trial. Use a permanent
marker and a 30 cm ruler to measure and mark a point on the bottom
of the box from the inside which is 3.5 cm away from the light
sensor opening. Subsequently, use a protractor to draw out a circle
which has a radius of 2.5 cm, using the marked point as the centre
of the circle.
3. Using sellotape, cover all sides of the box with a black,
opaque material, such as a bin bag. Make sure that the beaker
opening and light sensor opening are not covered by the bin bag
material.
4. Using a pair of scissors, cut out a 7 x 7 cm square out of
the bin bag material. Pierce a hole into the center of this square
using scissors.
5. Stick one side of this square to the top of the box in order
to cover the beaker opening. It is essential that this square is
only stuck on one side so that it acts as a flap which can be
opened and closed to reveal and cover the beaker opening.
B) PART 2: setting up the apparatus 6. Connect the LabQuest data
logger to the computer and, to it, connect the Vernier light
sensor. Make sure that the light
sensor is switched to collect data in the “0-600 lux” range,
given that this is the most suitable range for this experimental
procedure.
7. Make sure that the Logger Pro data logging software is loaded
and set the sampling duration to 60 seconds and the sampling rate
to 2 samples per second.
38 “Inverse Square Law Formula.” SoftSchools.com. Soft Schools.
Accessed November 24, 2019.
http://www.softschools.com/formulas/physics/inverse_square_law_formula/82/.
39 "Chemiluminescence with Luminol." North Carolina State
University Department of Chemistry. n.d. Raleigh. PDF. 7
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8. Stick the end of the light sensor into the light sensor
opening in the box, ensuring that the end of the light sensor
fits
snugly in the opening but doesn’t protrude into the box itself.
Critically, the end of the light sensor should be situated 1 cm
from the face of the beaker inside the box.
C) PART 3: preparing the solutions 9. To prepare the different
alkaline luminol solutions, place a 500 cm3 beaker on the
electronic weighing scale and tare the
scale. Hereafter, weigh 100 grams of distilled water in the
beaker. 10. Transfer the distilled water in the beaker to a 250 cm3
volumetric flask and place this flask on the magnetic stirrer.
Set
the magnetic stirrer to the maximum speed. 11. Using a 10 cm3
graduated pipette, measure out the 5 cm3 of NaOH. Add this NaOH to
the volumetric flask. 12. Using the electronic weighing scale,
weigh out the appropriate mass of luminol on a weigh boat as
indicated in Table 1.
Add this luminol to the volumetric flask as well and allow it to
dissolve. Once all the luminol has dissolved, top up the volumetric
flask with distilled water to the 250 cm3 mark.
13. Repeat steps 9 to 12 in order to prepare the 3 different
alkaline luminol solutions, as indicated in Table 1. 14. To prepare
the 3 different oxidizing solutions, repeat steps 9 to 13. However,
instead of adding NaOH and luminol,
measure out the appropriate volumes of H2O2 and mass of
K3[Fe(CN)6] respectively, as indicated by Table 2. D) PART 4:
conducting the experiment 15. Using a 10 cm3 volumetric pipette,
transfer 10 cm3 of an appropriate alkaline luminol solution into a
50 cm3 beaker.
Transfer this beaker through the beaker opening into the box and
place it on the indicated circular marking. Subsequently, close the
beaker opening flap.
16. Use a 10 cm3 syringe to withdraw 10 cm3 of an appropriate
oxidizing solution. Start recording the illuminance and, after 5
seconds, inject the oxidizing solution through the hole in the
beaker opening flap into the beaker containing the alkaline luminol
solution.
17. Allow the light sensor to record the illuminance for the
remaining 55 seconds of the sampling duration. Once the sampling is
complete, make sure to save the data file.
18. Repeat steps 15 to 17 two more times to achieve three
experimental trials for the run in question, making sure to create
a new data file for each trial.
19. Repeat steps 15 to 18 for the remaining runs of the
experiment, ensuring that three experimental trials are conducted
per run. Utilize Table 3 to determine the appropriate solutions to
use for each run.
E) PART 5: data processing 20. Open up a chosen data file on the
Logger Pro data logging software and, under the ‘Analyze’ tab,
select ‘Examine’. This
function will allow you to trace the graphed data points. 21.
Locate the spike in illuminance caused by the emission of light by
the reaction. Using the examine function, identify the
data point at the peak of the spike (the “final” illuminance)
and the data point prior to the spike (the “initial”
illuminance).
22. Using these two points, calculate the change in illuminance
and the change in time. The change in illuminance can be calculated
by subtracting the “final” illuminance value from the “initial”
illuminance value while the change in time can be calculated by
subtracting the “final” time value from the “initial” time value
(refer to Figure 7).
23. Repeat steps 20 to 22 for all 3 trials of each run of the
experiment in order to determine the change in illuminance and
change in time for each experimental trial.
Table 1: the different alkaline luminol solutions and their
contents Solution number Contents Concentration of luminol
dianions
Millimoles per unit volume / mmol dm-3
1 0.20 g luminol in 5 cm3 of 1 M NaOH and diluted to 250 cm3
with distilled water 6.80
2 0.10 g luminol in 5 cm3 of 1 M NaOH and diluted to 250 cm3
with distilled water 3.40
3 0.05 g luminol in 5 cm3 of 1 M NaOH and diluted to 250 cm3
with distilled water 1.70
A sample calculation for the concentration of luminol dianions
is shown below. As established previously, the concentration of
luminol dianions is equal to the concentration of luminol given
that 1 luminol molecule reacts with two OH- ions to form 1 luminol
dianion (as illustrated in Figure 2). Given this, the concentration
of luminol dianions and the concentration of luminol will be made
synonymous henceforth in this investigation.
Sample calculation for the concentration of luminol dianions in
solution 1 olar amount of luminol .00170 molesM = massmolar mass
=
0.200117.16 = 0
olar amount of luminol molar amount of luminol dianions, ∴ molar
amount of luminol dianions 0.00170 mol M = = oncentration of
luminol dianions .00680 mol dm .80 mmol dmC = volume
molar amount = 0.2500.00170 = 0 −3 = 6 −3
8
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Table 2: the different oxidizing solutions and their contents
Solution number Contents Concentration of O2 Millimoles per unit
volume / mmol dm-3
4 3.00 g K3[Fe(CN)6] in 10.0 cm3 of 5% H2O2 and diluted to 250
cm3 with distilled water 29.4
5 3.00 g K3[Fe(CN)6] in 5.0 cm3 of 5% H2O2 and diluted to 250
cm3 with distilled water 14.7
6 3.00 g K3[Fe(CN)6] in 2.5 cm3 of 5% H2O2 and diluted to 250
cm3 with distilled water 7.35
A sample calculation for the concentration of oxygen is shown
below. Evidently, the concentration of O2 is half the concentration
of H2O2 given that two H2O2 molecules decompose into one O2
molecule (as illustrated in the BACKGROUND INFORMATION
section).
Sample calculation for the concentration of oxygen in solution 4
5% H2O2 means that for every 100 cm3 of solution there is 5 grams
of H2O2 Therefore, 10 cm3 5% H2O2 will have 0.50 grams of H2O2
olar amount of H O .0147 molesM 2 2 =mass
molar mass =0.5034.01 = 0
mol of H O 1 mol of O , ∴ molar amount of O .00735 moles2 2 2 :
2 2 = 20.0147 = 0
oncentration of O .0294 mol dm 9.4 mmol dmc 2 = volumemolar
amount = 0.250
0.00735 = 0 −3 = 2 −3
Table 3: the relative concentrations of each reactant in the
runs of the experiment and the solutions used in each run
Run Concentration of luminol Concentration of O2 Alkaline
luminol solution used Oxidizing solution used
A Initial (6.80 mmol dm-3) Initial (0.294 mol dm-3) Solution 1
Solution 4
B Initial (6.80 mmol dm-3) Half (0.147 mol dm-3) Solution 1
Solution 5
C Initial (6.80 mmol dm-3) Quarter (0.00735 mol dm-3) Solution 1
Solution 6
D Half (3.40 mmol dm-3) Initial (0.294 mol dm-3) Solution 2
Solution 4
E Quarter (1.70 mmol dm-3) Initial (0.294 mol dm-3) Solution 3
Solution 4
DIAGRAM OF APPARATUS
Diagram 1: a diagram of the light-blocking box with the
dimensions of the box, the beaker opening and the light sensor
opening - self-made diagram
Diagram 2: a cross-sectional diagram of the apparatus setup
showing the contents of the light-blocking box - self-made
diagram
SAFETY AND ENVIRONMENTAL CONSIDERATIONS In order to ensure the
safety of the experimenter, appropriate lab safety attire such as
safety goggles and lab coats should be worn throughout the
experiment. This is especially important when handling the
chemicals as they are all harmful in case of skin contact, eye
contact and inhalation , , , . Experimenters must also be aware of
the location of all emergency facilities 40 41 42 43
40 "SAFETY DATA SHEET Luminol." Cayman Chemical, April 21, 2019.
PDF. 9
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in the near vicinity, including the eye-wash station, shower and
fire extinguisher, before conducting the experiment. In case of
skin or eye contact with any of the chemicals, make sure to take
off any contaminated clothing and rinse the skin/eyes with
water.
It is necessary to dispose of the reactant solutions in a
suitable chemical disposal container given that some of the
reactants, particularly potassium ferricyanide and sodium
hydroxide, are toxic to marine life , . A further environmental 44
45consideration to note is the need to make larger batches of the
reactant solutions than necessary, meaning that a substantial
volume of the solutions will be unused and disposed of. This is due
to the fact that very small masses of the reactants are used to
make these solutions, meaning that the smallest batch size that
could be made for each solution using the available measuring
apparatus is 250 cm3. Given this, and the fact that only 10 cm3 of
each solution (alkaline luminol and oxidizing solution) is used per
trial, a total of 1500 cm3 of solution will be prepared while only
300 cm3 of solution will be used in the experiment.
DATA COLLECTION AND PROCESSING After conducting all trials and
runs of the experiment, the values for the initial illuminance and
final illuminance were recorded in Table 4 below:
Table 4: raw data collection showing the initial and final
illuminance and time for each trial of each run
Run
Concentration of luminol
moles per unit volume / mol
dm-3
Concentration of O2
moles per unit volume / mol
dm-3
Initial illuminance illuminance / lx
±0.0001 lx
Final illuminance illuminance / lx
±0.0001 lx
Initial time seconds / s ±0.0001 s
Final time seconds / s ±0.0001 s
Trial 1 Trial 2 Trial 3 Trial 1 Trial 2 Trial 3 Trial 1 Trial 2
Trial 3 Trial 1 Trial 2 Trial 3
A 0.0068 0.294 1.1036 2.3680 1.9269 9.8010 9.4908 8.7438 0.0750
0.0500 0.0875 0.1000 0.0750 0.1125
B 0.0068 0.147 1.6919 1.7506 1.7271 7.2733 7.1887 6.9556 0.0875
0.0750 0.0250 0.1250 0.1125 0.0800
C 0.0068 0.0735 3.1723 1.6684 1.6331 4.4735 3.0365 3.9375 0.0750
0.0875 0.1250 0.1000 0.1125 0.1750
D 0.0034 0.294 1.7506 1.9504 1.5744 5.1757 5.5582 3.4279 0.0750
0.0625 0.0750 0.1000 0.0875 0.0875
E 0.0017 0.294 1.9504 1.5509 2.3506 5.0945 4.6125 3.9984 0.0875
0.1625 0.0750 0.1325 0.2125 0.1000
Following this, it was necessary to calculate the changes in
illuminance and time for each trial of each run. Sample
calculations for the changes in illuminance and time and their
respective absolute uncertainties are shown below. Given that the
calculations for the change in illuminance and time only involved a
subtraction, the absolute uncertainties of the initial and final
illuminance and time were added in order to find the absolute
uncertainty of the change in illuminance and time respectively.
Critically, the absolute uncertainty for all of the change in
illuminance and change in time values is ±0.0002 lx and ±0.0002 s
respectively, given that all the illuminance and time values in
Table 4 have an absolute uncertainty of ±0.0001 lx / ±0.001 s.
Sample calculation for the change in illuminance for trial 1 of
run A hange in illuminance inal illuminance initial illuminance
.8018 .1036 .6974 lx (correct to 4 decimal points) c = f − = 9 − 1
= 8
Sample calculation for the change in time for trial 1 of run
A
hange in time inal time initial time .1000 .0750 .0250 s
(correct to 4 decimal points) c = f − = 0 − 0 = 0 Sample
calculation for the absolute uncertainty of the change in
illuminance/time
a b cΔ = Δ + Δ a .0001 .0001 = Δ = 0 + 0 a .0002 lx / 0.0002 s =
Δ = 0
a bsolute uncertainty of the change in illuminance/timeΔ = a b
bsolute uncertainty of the initial illuminance/timeΔ = a c bsolute
uncertainty of the f inal illuminance/timeΔ = a
The values for the change in illuminance and change in time for
each trial of each run were recorded in Table 5 below. Using these
values, the rate of reaction for each trial, the mean rate of
reaction and the absolute uncertainty of the mean rate of reaction
could be calculated.
41 "SAFETY DATA SHEET POTASSIUM FERRICYANIDE." Sigma Aldrich,
January 6, 2019. PDF. 42 "SAFETY DATA SHEET Hydrogen Peroxide 10%."
PeroxyChem, April 8, 2015. PDF. 43 "Sodium Hydroxide Safety Data
Sheet." LabChem, March 26, 2012. PDF. 44 Ibid. 45 "SAFETY DATA
SHEET POTASSIUM FERRICYANIDE." Sigma Aldrich, January 6, 2019.
PDF.
10
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Table 5: processed data showing the change in illuminance,
change in time and the rate of reaction for each trial as well as
the mean rate of reaction and its absolute uncertainty for each
run
Run
Concentration of luminol
moles per unit volume / mol
dm-3
Concentration of O2
moles per unit volume / mol
dm-3
Change in illuminance illuminance / lx
±0.0002 lx
Change in time seconds / s ±0.0002 s
Rate of reaction illuminance over time / lx s-1
Mean rate of reaction
illuminance over time / lx s-1
Absolute uncertainty of
the mean rate of reaction
illuminance over time / lx s-1 Trial 1 Trial 2 Trial 3 Trial 1
Trial 2 Trial 3 Trial 1 Trial 2 Trial 3
A 0.0068 0.294 8.6974 7.1228 6.8169 1.5000 1.5000 1.5000 5.7983
4.7485 4.5446 4.6466 0.1020
B 0.0068 0.147 5.5814 5.4381 5.2284 2.2500 2.2500 3.2500 2.4806
2.4169 1.5844 2.4488 0.0319
C 0.0068 0.0735 1.3012 1.3681 2.3044 1.5000 1.5000 3.0000 0.9375
0.9121 0.8881 0.9149 0.0720
D 0.0034 0.294 3.4250 3.6078 1.8535 1.5000 1.5000 0.7500 2.2834
2.4052 2.4713 2.3866 0.0940
E 0.0017 0.294 3.1441 3.0616 1.6478 2.7000 3.0000 1.5000 1.1645
1.0205 1.0985 1.0945 0.0720
Sample calculations for the rate of reaction in each trial, the
mean rate of reaction and the absolute uncertainty of the mean rate
of reaction are shown below.
Sample calculation for the rate of reaction for trial 1 of run B
ate .4806r = change in time
change in illuminance = 2.25005.5814 = 2 x s (correct to 5
signif icant f igures) l −1
Sample calculation for the mean rate of reaction for run B
(accounted for anomaly)
ean .4488M = number of trialssum of individual trials = 2
2.4806+2.4169 = 2 x s (correct to 5 signif icant f igures) l −1
Sample calculation for the absolute uncertainty of the mean rate of
reaction for run B:
bsolute uncertainty 0.0319 lx s (correct to 4 decimal places) a
= 2range = 2
max. value − min. value = 22.4806−2.4169 = ± −1
In Table 5 it can be seen that two rate of reaction values in
runs A and B are highlighted in orange - identifying them as
anomalies. This is evidenced by the fact that the removal of these
two values (5.7983 and 1.5844 lx s-1) substantially decreased the
absolute uncertainty of the mean rate of reaction in their
respective runs. For example, the removal of the 5.7983 lx s-1 data
point from run A reduced the absolute uncertainty of the mean rate
of reaction from ± 0.6269 lx s-1 to ± 0.1020 lx s-1. Ultimately,
these anomalous results were ignored when calculating the mean rate
of reaction in order to enhance the accuracy of this
calculation.
Additionally, qualitative observations were made throughout the
experiment and recorded in Table 6. These observations may aid in
the interpretation of the experiment's results. Table 6:
qualitative observations made throughout the experiment
Run(s) Qualitative observations
A, D, E In runs A, D and E, which used solutions with the
highest oxygen concentrations, light effervescence of oxygen
bubbles was observed in the oxidizing solution prior to the
reaction. This effervescence was most prominent when the oxidized
solution was disturbed (e.g. when a syringe was used to withdraw
the solution). While the oxygen effervescence was also observed in
oxidizing solutions with lower concentrations of oxygen (e.g. runs
B and C), this was much fainter than in runs A, D and E.
All runs Luminol was least soluble in the runs where it was used
at the highest concentration (runs A, B and C), yet was also
observed to be relatively insoluble in the other runs of the
experiment. Due to this limited solubility, the alkaline luminol
solutions needed to be stirred excessively (particularly in runs A,
B and C) in order to dissolve as much luminol as possible. Despite
these efforts, a precipitate of undissolved luminol remained at the
bottom of the alkaline luminol solutions used (a larger amount of
precipitate was observed in runs A, B and C), indicating that not
all of it had dissolved.
All runs In all of the trials it was seen that the baseline
illuminance reading fluctuated slightly above zero (usually at 0.7
lux).
Ultimately, the mean rate of reaction determined in Table 5 was
graphed against the concentration of luminol and O2, as is shown in
Graphs 1 and 2. In these graphs, the error bars represent the
absolute uncertainty of the mean rate of reaction.
11
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Graph 1: concentration against mean rate of reaction graph for
luminol
Graph 2: concentration against mean rate of reaction graph for
oxygen
PERCENTAGE ERROR Given that no experimental value was determined
in this investigation, a percentage error could not be
established.
ANALYSIS AND CONCLUSION The aim of this experiment was to
experimentally derive the rate law for the chemiluminescent
oxidation of luminol. As was already evident in Table 5, halving
the concentration of either reactants resulted in the mean rate of
reaction approximately halving as well. For example, the ratio of
the mean rates of reaction in runs A and D (where run D utilizes
half the concentration of luminol used in run A) is approximately 2
( ), illustrating how halving the concentration of .95 2.3866
4.6466 = 1 ≈ 2 luminol also halved the mean rate of the
reaction. This trend is also seen in Graphs 1 and 2, where a linear
relationship was illustrated between the concentrations of both
luminol and O2 and the mean rate of reaction. Critically, this
relationship suggests that the reaction is first-order with respect
to both of these reactants, and hence the experimentally derived
rate law is .ate [luminol][O ]r = k 2
These findings don’t only support the steady-state approximation
which informed my initial hypothesis of the rate law, but also
illustrate the effects of varying concentrations on the reaction
rate, as was discussed in the BACKGROUND INFORMATION section. As
the concentration of either reactants was increased, the frequency
of collisions between reactant particles also increased, thus
resulting in an increased rate of photon emission and hence an
increased rate of reaction. Given that the reaction was first order
with respect to both luminol and O2, halving the concentration of
either reactant halved the rate of the reaction. This is why the
mean rate of reaction shown in Table 5 for runs B and D (which
utilized half the initial concentration of O2 and luminol
respectively) and runs C and E (which utilized a quarter of the
initial concentration of O2 and luminol respectively) are so
similar.
Furthermore, these findings also confirm that the rate
determining step in this reaction mechanism is step 3. Prior to
conducting the experiment, there was some uncertainty as to which
reactants were part of the rate law given that it wasn’t known
which step is the rate-determining step in the mechanism. However,
given that this experiment has illustrated that O2 is part of the
rate law, step 3 must be the rate-determining step since it
involves the reaction between the luminol dianions and O2.
Critically, given that the formation of the luminol dianions is
dependent on the reaction between luminol and the OH
- ions (steps 1 and 2), both luminol and OH- ions must also be
included in the rate law. While the first-order nature of luminol,
as was hypothesized by the steady-state approximation, was
confirmed experimentally, the effects of the concentration of OH-
ions wasn’t investigated, the reasons for which were explained
prior to the experiment.
EVALUATION
Evidently, the results of the experiment don’t fully answer my
research question; “What is the rate law of the chemiluminescent
oxidation of luminol (C8H7N3O2 ) by sodium hydroxide (NaOH ) and
hydrogen peroxide (H2O2 )?”, given that the experiment didn’t
consider the effects of the concentration of OH- ions on the rate
of the reaction and I've therefore not been able to construct a
rate law for the reaction including all of its reactants. On one
hand, it could be argued that the fact that the orders of reaction
with respect to both luminol and O2 have been supported by the
steady-state approximation strengthens the case that the
hypothesized reaction order with respect to OH- ions (first-order)
is accurate. However, another interpretation may suggest that,
given that the formation of the luminol dianions depends on the
reaction between one
12
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luminol molecule and two OH- ions, the reaction should be second
order with respect to the OH- ions. Evidently, it is difficult to
conclude which of these orders of reaction is experimentally
observed.
The high precision of the collected data is also apparent when
considering the uncertainty of my results. This is demonstrated by,
for example, the high coefficient of determination (R2) values in
Graphs 1 and 2 (R2 = 0.9986 and 0.9901 respectively), which
indicate that there is very minimal variance of the data points
from the linear trendline . However, it 46can further be seen that
the coefficient of determination value is lower for Graph 3 than
for Graph 2 , suggesting that a random error affected the precision
of the data collected for the varying concentrations of O2. A
possible source of this error was a random error commented upon in
the qualitative observations in Table 6; the effervescence of the
oxidizing solution. The fact that hydrogen peroxide decomposes into
oxygen gas was overlooked in this investigation and resulted in the
inevitable loss of oxygen from the oxidizing solution throughout
the experiment. The effervescence of oxygen was evident in all
solutions yet was most prevalent in the oxidizing solution used in
trials A, D and E which contained the highest concentration of O2.
This observation, alongside the fact that the runs and trials in
this investigation weren't conducted in chronological order,
would’ve most likely resulted in greater variance in the collected
data as the concentration of O2 in the oxidizing solution would
decrease over time as O2 was lost as gas. This variance was also
reflected in the small error margin of the data. The absolute
uncertainty of the mean rate of reaction values was very small,
ranging from ±0.0319 lx s-1 to ±0.1020 lx s-1, and the error bars
in Graph 1 and 2 are thus barely visible, supporting the high
precision of the data. However, while the trendline in Graph 1
passes through all of these error bars, the trendline in Graph 2
does not, which demonstrates the relatively lower precision of the
collected data in Graph 2 in comparison to the data in Graph 1.
This observation could also be explained by the oxygen loss
observed and may further suggest that the error margins for this
investigation were too small (and hence underestimated).
One way in which the accuracy of the collected data was ensured
was through the identification and removal of anomalous data in
Table 5, effectively reducing the likelihood of an erroneous
correlation being identified. Despite this, both Graphs 1 and 2
demonstrated a systematic shift of the trendline, as is indicated
by the extrapolated trendline which intersects the y-axis below the
origin. The cause of this systematic shift may partly be due to a
prominent systematic error which was commented upon in the
qualitative observations in Table 6; the solubility of luminol in
the alkaline luminol solution. This phenomenon was observed
throughout the experiment, as the beakers with the alkaline luminol
solutions had a layer of undissolved luminol precipitate despite
intensive and sustained stirring. Due to this, the concentration of
luminol was consistently underestimated, given that the actual
concentration of luminol in the alkaline luminol solutions was
consistently lower than was theoretically calculated, which
explains the systematic (rightward) shift of the trendline. Prior
to conducting the experiment, no consideration was given to the
solubility of luminol in water; the primary solvent used.
Unsurprisingly, subsequent research revealed that luminol is
“comparatively insoluble in water, but quite soluble in base”, thus
explaining the insolubility of luminol in water and further
indicating that the pH of the alkaline luminol solution (despite
the addition of the NaOH base) was too low. Another possible cause
of this systematic shift would be due to the underestimation of the
change in illuminance caused by the emission of photons. In the
qualitative observations in Table 6, it was noted that the baseline
illuminance value recorded in all trials of all runs fluctuated
slightly above 0 lx. Due to this zero error, the change in
illuminance would’ve been consistently underestimated throughout
the experiment, meaning that the rates of reaction would be
consistently underestimated (given that the rate of reaction was
calculated as the change in illuminance over time). This,
ultimately, provides an explanation for the systematic (downwards)
shift of the trendline. Both of these systematic errors - the
solubility of luminol and underestimation of the change in
illuminance - are valid explanations for the systematic shift of
the data, and aiming to overcome both of these errors would
therefore improve the accuracy of the data collection.
ERRORS AND IMPROVEMENTS SYSTEMATIC ERRORS 1. Solubility of
luminol: the solubility of luminol was a significant error in this
investigation as it was thought to partly
cause the systematic shift of the data, as explained previously.
One way to increase the solubility of luminol, and thus overcome
this error, would be to adjust the contents of the alkaline luminol
solutions to initially contain less distilled water or more NaOH,
both of which will increase the concentration of OH- ions in the
solution and hence increase the solution’s pH. Alternatively, it
would be possible to add other bases that wouldn’t react with
either NaOH or luminol (e.g. sodium carbonate and ammonium
carbonate) in order to raise the pH of the alkaline luminol
solution and thus increase the solubility of luminol.
2. Zero error: the other potential cause of the systematic shift
of the data was due to the underestimation of the change in
illuminance, caused by the positive zero error in the light sensor.
This zero error indicates that the inside of the light-blocking box
wasn’t entirely dark and that there were therefore sources of
external light whose illuminance was recorded from inside the box.
This error is significant given that it may have caused the
systematic shift of the data, but it doesn’t refute the trendline
found given that the same apparatus (light-blocking box) was used
throughout the experiment. To improve upon this error, one could
construct a light-blocking box from a more suitable,
light-blocking
46 Frost, Jim. 2018. How To Interpret R-squared in Regression
Analysis. February 24. Accessed November 12, 2019.
https://statisticsbyjim.com/regression/interpret-r-squared-regression/.
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material (e.g. a large film canister) in order to better block
out external light. Furthermore, the entire experiment could be
conducted in a dark laboratory room in order to reduce the amount
of external light which needs to be blocked by the light-blocking
box.
3. Limited range and number of reactant concentrations: this
experiment tested only 3 different concentrations of each reactant
in order to determine their respective orders of reaction. While
this resulted in an observable trend, the limited range and number
of different concentrations meant that it was difficult to deduce
if, for whatever reason, the mean rate of reaction in one (or more)
runs was anomalous, given that a trendline needed to be fitted to
only 3 different data points in each graph. Ultimately, this error
can be overcome by using a greater number and range of different
reactant concentrations, effectively increasing the reliability of
the trends observed in Graphs 1 and 2. For example, a wider range
of 5 different concentrations could be investigated.
RANDOM ERRORS 1. Loss of oxygen from the oxidizing solution: the
loss of oxygen from the oxidizing solution is a significant error
as it
was thought to greatly increase the variance and thus reduce the
precision of the collected data. Critically, this error was caused
by the flawed assumption that the concentration of O2 in the
oxidizing solutions would be equal to the concentration of H2O2
since the loss of oxygen gas from the oxidizing solution was
overlooked. Ultimately, this error could be improved upon by
creating the oxidizing solutions such that their O2 concentration
is lower, given the fact that the oxidizing solutions with a higher
O2 concentration were seen to produce the most effervescence. This
lower concentration could be achieved by using a lower percentage
concentration of H2O2 or by adding lower volumes of H2O2 to the
oxidizing solutions. While this methodological improvement won’t
eliminate the presence of effervescence it will greatly reduce it
such that the assumption that the concentration of O2 in the
oxidizing solutions would be equal to the concentration of H2O2 is
more accurate.
2. Limited number of trials: in the experiment, only 3 trials
for each run was conducted. While this allowed me to calculate the
mean rate of reaction using three different data points in most
runs, removing an anomalous data point (as was done in runs A and
B) would mean that the mean of only two values was taken, which
isn’t very reliable. Moreover, given that the absolute uncertainty
of the mean was calculated as the range of data points divided by
2, removing anomalous data points would possibly result in the
underestimation of the error margins in the experiment, as was
outlined in Graph 2. Ultimately, in order to improve upon this
error one would need to conduct a larger number of trials, say
five, in the experiment.
EXTENSIONS An obvious extension to this investigation would be
to determine the order of the reaction with respect to OH- ions,
but the methodology required to do this is beyond my level of
understanding and requires the use of more elegant equipment due to
the presence of three different reactants. A more realistic
extension to this investigation would be to consider the energy
lost as heat in the reaction. One of the assumptions made in this
experiment is that the energy released by the reaction is released
as light, following the definition of “cold light”. However, this
is a slight simplification given that there will always be a
certain amount of energy which is released as thermal energy,
although this amount of energy may not be so significant in this
reaction. Nevertheless, by keeping track of the temperature change
of the reaction mixture as the reaction progresses, one could take
into account the energy lost as heat and integrate this with the
energy lost in the form of light in order to get a more accurate
estimation of the enthalpy change of the reaction. This could,
ultimately, be used to derive a more accurate value for the change
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