Final Photocatalysis Lab Report (1) (1)

March 6th
, 2014
Professor LiangFang Zhang
University of California San Diego
9500 Gilman Dr.
La Jolla, CA 92093
Dear Professor LiangFang Zhang,
This report is an analysis of the methylene blue degradation kinetic constant as the amount of
hydrogen peroxide is varied. By increasing the amount of hydrogen peroxide, it should be
possible to increase the production of hydroxyl radicals, which should lead to faster methylene
blue degradation kinetics when exposed to UV light. The highest methylene blue degradation
kinetic constant was determined to be approximately 0.12 min-1
, which corresponded to the
highest added hydrogen peroxide volume of 10 mL (a concentration of 36 mM). If there are any
questions, please contact the group.
Sincerely,
Group 3
Felipe Alcantara Sam Chan
Henry Hsieh Vincent Li

Analysis of methylene blue degradation kinetic constant as the
amount of hydrogen peroxide is varied
March 6th
, 2014
Group Member Section
Felipe Alcantara
Technical Memo, References and Citations, Figures and Tables, Briefing
Presentation
Sam Chan Theory and Background
Henry Hsieh Results and Discussion
Vincent Li Letter of Transmittal, Abstract, Introduction, Methods, and Conclusions

3
Abstract:
The goal of this experiment was to study the effect of hydrogen peroxide (H2O2) on the
methylene blue degradation kinetic constant. Specifically, 5mL and 10mL (correspondingly a
concentration of 18 mM and 36 mM) of H2O2 were tested. With increasing volume of H2O2, it
was expected that the methylene blue degradation kinetic constant first increases, and then
decreases. The resultant kinetic constants from the added H2O2 practically coincided with the
expected trend. It was observed the resultant kinetic constant was respectively 0.07 min-1
and
0.12 min-1
for 5mL and 10mL of added H2O2. As a result, increasing the concentration of H2O2
can lead to faster methylene blue degradation. However, due to insufficient kinetic study for
higher H2O2 concentration, the decreasing of the methylene blue degradation kinetic constant
from excess H2O2 was not observed.
Introduction:
Wastewaters from industries, textile factories, and laboratories contain dyes that are
harmful to humans due to its endocrine disrupting properties, and it has been noted that as much
as 1–20% of the total world production of dyes is released to the environment. These pollutant
dyes can further generate other toxic byproducts through oxidation and can also prevent light
from entering into water, which poses a significant threat to all aquatic life and to the
ecosystem.1
As a result, it is imperative to find an effective method to degrade textile dye.
Various dye degradation methods, such as adsorption on activated carbon, ultrafiltration, and
reverse osmosis, has been studied.1,2
However, these methods only generate a secondary
pollutant as the dye is transferred from wastewater to another solvent, which then requires a
secondary treatment. Recently, attention has been given to treating textile dyes through

4
photocatalysis due to its low cost and effectiveness in converting harmful dyes into harmless
chemicals such as CO2, HCl, and water.3,4
During the photocatalysis reaction, TiO2/dye mixture pH, precursor TiO2 pH, type and
amount of catalysis, dopant content, temperature, and oxidizing agents can affect the dye
degradation kinetics.1
When studying the effect of oxidizing agents (such as H2O2) on the
photocatalytic degradation of dyes, it was observed that increasing H2O2 led to an increase in
kinetics because the radicals generated (such as hydroxyl and various other radicals) can serve as
both effective oxidizing and reducing agents, which greatly limits the undesired electron-hole
recombination at the TiO2 surface.5,6
From Huang et al. and Konstantinou et al. it is also
interesting to note that there should be an optimal amount of H2O2 added to obtain the highest
dye degradation kinetic.5,7
This is because catalytic activity can actually decrease due to radical-
radical recombination and H2O2 adsorption onto TiO2 surface if there is an excess of H2O2. As a
result, the overall goal of this experiment was to study the methylene blue degradation kinetic
constant as the amount of H2O2 was varied.
Theory and Background:
Photocatalysis involves the use of UV light, catalysts, and oxidizers to degrade complex
organic chemicals. Typical catalysts used in photocatalysis are semiconductor particles such as
TiO2, ZnO, and Fe2O3. TiO2 is generally used due to its stability, biocompatibility, optical and
electrical properties. As UV light illuminates the TiO2 particles, electrons in the valence band
can be promoted to the conduction band, generating an electron-hole pair (Reaction 1). The
overall goal of the photocatalytic reaction is to prevent the electron-hole pair recombination by
having the electrons and holes reacting with other molecules present in the system. By

5
preventing the recombination, the dye can also be degraded by the generated electrons and holes
(Reactions 2 and 3). However, the electrons can also react with oxygen molecules to produce
superoxide anions (Reaction 4). Similarly, the holes can react with water and hydroxide anions to
generate hydroxyl radicals (Reactions 5 and 6). These reactions do not hinder dye degradation
due to the fact that the reactive radicals are likely to take electrons from the dye molecules,
which destabilizes them for degradation (Reaction 7). Furthermore, an oxidizing agent such as
H2O2 can be introduced to increase the amount of hydroxyl radicals in the system (Reactions 8-
10), which should lead to higher dye degradation due to Reaction 7.1,8
(1)
(2)
(3)
( ) (4)
( ) (5)
( ) (6)
(7)
(8)
(9)
( ) (10)
Although H2O2 can facilitate dye degradation, an excess of H2O2 can actually lower the
dye degradation rate because of H2O2 adsorption onto TiO2 surfaces, TiO2 holes depletion, and
hydroxyl radical consumption (Reactions 11-14).1
( ) (11)
(12)

6
(13)
(14)
To study the effect of H2O2 on the dye degradation kinetics, Langmuir-Hinshelwood (L-
H) kinetics was applied (Equation 1),7
(1)
where r is the reaction rate in ppm/min, C is the concentration in ppm, t is time in min, kLH is the
L-H rate constant in ppm/min, and KL is the Langmuir adsorption constant in ppm-1
.
At low concentrations, 1+KLC in Equation 1 can be approximated as 1, and the kLH and
KL terms can be combined into an apparent rate constant, k, with units of min-1
. Then, the
differential dye degradation rate law takes the form of Equation 2. 7,9
(2)
Due to the fact that the photocatalytic reactor system has a UV-lamp reactor tank and a
holding tank instead of one single reaction tank, the left-hand side of Equation 2 was multiplied
by the total liquid volume in the system (VT) with unit of mL, and the right-hand side was
multiplied by the liquid volume inside the UV-lamp reactor tank (VR) with unit of mL to obtain
Equation 3.
(3)
Then, integrating Equation 3 results in Equation 4 (see Appendix for the complete
derivation of Equation 4).
( ) (4)
Equation 4 was based on the assumption that VT remained constant throughout the entire
kinetic study. However, TiO2/methylene blue mixture aliquots were taken from the

7
photocatalytic reactor system during the experiments, so the total volume was not constant. As a
result, VT in Equation 3 was approximated by Equation 5,
(5)
where V0 is the initial mixture volume in mL, α is an approximated average volumetric rate of
mixture taken from the system during aliquot sampling. In this experiment, approximately 4 mL
of mixture were taken from the system every 2 minutes, which resulted in an α of 2 mL/min.
Substituting Equation 5 into 3 and integrating, Equation 6 was obtained (see Appendix
for the complete derivation of Equation 6).
( ) ( ) (6)
Using an UV-Vis spectrophotometer, absorbance were measured at each corresponding
sampling time (Tables 2 and 3 in Appendix). To study the mixture concentration progression as a
function of time (Equations 4 and 6), it was necessary to convert the measured absorbance into
concentration. As a result, Beer’s law (Equation 7) was used to obtain a calibration curve,10
(7)
where ε is the absorption coefficient of methylene blue in cm-1
ppm-1
, l is the path length (1 cm),
and c is the concentration in ppm.
Methods:
10 mg of methylene blue was dissolved in 1000 mL of DI water to generate an
approximately 10 ppm stock solution. Then, 250 mL of the stock solution was dissolved in
another 250 mL of DI water to form an approximately 5 ppm solution before UV-Vis
spectrophotometry was performed using the SpectroVis Plus spectrophotometer and the Logger
Pro software. A full wavelength scan (between 400 nm and 720 nm) of the 5 ppm methylene

8
blue solution was measured to obtain the wavelength corresponding to the highest absorbance
value (Figure 7 in Appendix). Then, all other absorbance measurements were taken at this
wavelength. Next, the methylene blue solution was diluted in a range of 0.1 to 0.9 times the
original 5 ppm solution (labeled as micro-cuvette 1 to 9 in Figure 1).
Figure 1. Methylene blue diluted in increments of 0.1 between 0.1 to 0.9 times (labeled as 1 to 9) the original 5 ppm
solution (labeled as 10).
The corresponding absorbance was measured, and a Beer’s law calibration curve was
generated. Next, 0.9 g of TiO2 and 500 mL of DI water were placed into the holding tank. After
the TiO2 solution was well mixed with a stir bar stirring at about 175 rpm, a Masterflex
peristaltic pump was utilized to circulate the TiO2 solution through the entire photocatalytic
reactor system. This ensured a homogeneous deposition of TiO2 inside the 350 mL 9 Watt UV-
lamp reactor. Next, 200 mL of the 10 ppm methylene stock solution was added to the holding
tank, and the solution was circulated around the system for 3 minutes to ensure a homogenous
TiO2/methylene blue mixture throughout the entire photocatalytic reactor system (Figure 2).

9
Figure 2. Photocatalytic reactor system with holding tank, peristaltic pump, and a 350 mL 9 watt UV-lamp reactor.
The degrading methylene blue is circulating clockwise with flow rates and magnetic stir speed held constant during
both experimental trials.
After that, the absorbance of the initial TiO2/methylene blue mixture was measured
before 5 mL (first trial), and 10 mL (second trial) of a 6% by volume H2O2 were added to the
holding tank. The UV-lamp was then immediately turned on, and roughly 4 mL aliquots of
TiO2/methylene blue mixture were taken from the holding tank every two minutes. The aliquots
were filtered (to remove the TiO2 particles) into a micro-cuvette, and the absorbance of the
degrading methylene blue at each sampling time were measured until the absorbance of the
methylene blue reached below 0.05.
Results and Discussion:
After a full wavelength scan between 400 nm and 720 nm for the 5 ppm methylene blue
solution was done, the maximum absorbance of 0.996 was observed at 659 nm (Figure 7 in
Appendix). After the wavelength corresponding to the maximum absorbance was obtained, all
absorbance were then measured while holding that wavelength constant. By taking absorbance
measurement at the wavelength corresponding to the highest absorbance, the most precise
absorbance readings can be obtained as absorbance fluctuation errors can be minimized. The

10
absorbance for the range of diluted methylene blue solution were used to construct a Beer’s law
calibration curve (Table 1 in Appendix and Figure 3), which made it possible to obtain
concentration from absorbance.
Figure 3. Calibration curve of average absorbance versus concentration (ppm), and it shows a methylene blue
absorption coefficient of 0.19 cm-1
ppm-1
.
A first order L-H kinetic rate plot was constructed, which assumed the total mixture
volume was unchanged during the kinetic study experiment (Figure 4). From Equation 4, the
slope in Figure 4 represented the methylene blue degradation kinetic constant. The degradation
kinetic constants for 5 mL and 10 mL (18 mM and 36 mM) were observed to be respectively
0.07 min-1
and 0.12 min-1
.
y = 0.1947x
0
0.2
0.4
0.6
0.8
1
1.2
0 1 2 3 4 5 6
AverageAbsorbance
Concentration (ppm)

11
Figure 4. First order L-H kinetic rate plot of (VT/VR)*ln(C/C0) versus time constructed from equation 4. The
degradation kinetic constants for 5 mL and 10 mL were respectively 0.07 min-1
and 0.12 min-1
However, it was noted that the total volume of TiO2/methylene blue mixture was
decreasing as aliquots were taken out of the photocatalytic system. To account for the decreasing
total volume, the relationship between ( /VR)*ln(C/C0) and ln(V0/(V0- t)) (equation 6) was
plotted (Figure 5) to obtain a more realistic methylene blue degradation kinetic constant.
Although the total volume would be better modeled as a decreasing step function due to aliquots
were only taken at sampling time, the total volume was assumed to decrease in a linearly
continuous fashion (a constant extraction rate) so that the integration of Equation 3 was possible.
From Figure 5, the resultant methylene blue degradation kinetic constants were respectively 0.06
min-1
and 0.12 min-1
for added H2O2 volume of 5 mL and 10 mL.

12
Figure 5. First order L-H kinetic rate plot of ( /VR)*ln(C/C0) with unit of min-1
versus ln(V0/(V0- t)) curve
constructed from equation 6. The degradation kinetic constants for 5 mL and 10 mL were respectively 0.06 min-1
and 0.12 min-1
Since kinetic constants from both Figure 4 and Figure 5 were similar, it is arguable that
the decreasing total volume from aliquots taken from the system did not significantly affect the
methylene blue degradation kinetic constant. As a result, the final degradation kinetic constants
were obtained from the plot that better resembles the L-H kinetics (Figure 4), which
corresponded to a degradation kinetic constant of 0.07 min-1
and 0.12 min-1
for 5 mL and 10 mL
of added H2O2, respectively. From these resultant degradation kinetic constants, it was observed
that the degradation kinetic constant did increase with increasing H2O2 concentration. However,
decreasing degradation kinetic constant from excess H2O2 was not observed due to insufficient
kinetic studies with higher H2O2 concentration.
Lastly, it was observed that both Figure 4 and Figure 5 displayed an inflection-like curve
instead of a linear straight line as predicted by the L-H kinetics. This inflection-like behavior can

13
be seen as three regions: the initial, the middle, and the end region (Figure 6). When the H2O2
was introduced initially, the sudden increase in H2O2 may have led to a large degradation
kinetics due to rapid hydroxyl radicals formation inside the UV-lamp reactor (initial region).
After the H2O2 and hydroxyl radicals become equilibrated throughout the photocatalytic system,
the degradation kinetics stabilizes (middle region). Finally, the overall degradation kinetics
towards the end of the experiment increased again most likely because the relative concentration
of H2O2 to concentration of methylene blue is greatly increased (end region). This can be seen
due to a higher H2O2 concentration from decreasing total mixture volume as aliquots were
extracted, and a lower methylene blue concentration as the methylene blue is degraded. As a
result, it is possible that the degradation kinetic increases as the overall exposure time of the
TiO2/methylene blue mixture is extensively prolonged.
Figure 6. Three regions of an inflection-like curve: the initial, the middle, and the end region.

14
Conclusion:
The purpose of this experiment was to study the effect of H2O2 on the degradation of
methylene blue using UV-light. Based on the data collected, it was observed that an increasing
amount of H2O2 led to a higher methylene blue degradation kinetic constant (faster degradation).
It was determined that the methylene blue degradation kinetic constants for adding 5 mL and 10
mL (corresponding concentration of 18 mM and 36 mM) of H2O2 were respectively 0.07 min-1
and 0.12 min-1
. This followed the expected trend of increasing H2O2 would lead to a higher
methylene blue degradation kinetic constant due to increase in hydroxyl radicals (Reaction 7-10).
However, it was observed that the range of H2O2 concentration tested may not contain the
optimal concentration of H2O2. To find the optimal concentration of H2O2, higher H2O2
concentration should be tested until the methylene blue degradation kinetic constant decreases. It
was also noted that the degradation kinetic constants were only an approximation due to the total
volume of the TiO2/methylene blue mixture was decreasing as aliquots were taken out of the
photocatalytic reactor system. To obtain more accurate methylene blue degradation kinetic
constant for future experiments, the measured aliquots should be replaced into the holding tank.
This should allow kinetic study experiments to be done without significantly changing the total
TiO2/methylene blue mixture volume within the photocatalytic reactor system.
References:
1. Akpan, U.G.; Hameed, B. H. J. Hazard. Mater. 2009, 170, 520-529.
2. Tang, W. Z.; An, H. Chemosphere. 1995, 9, 4157-4170.
3. Aramendia, M. A.; Marinas, A.; Marinas, J. M.; Moreno, J. M.; Urbano, F. J. Catal.
Today. 2005, 101, 187-193.

15
4. Gaya, U.I.; Abdullah, A.H. J. Photochem. Photobiol., C. 2008, 9, 1-12
5. Huang M.; Xu, C.; Wu, Z.; Huang, Y.; Lin, J.; Wu, J. Dyes Pigm. 2008, 77, 327-334.
6. Carp, O.; Huisman, C. L.; Reller, A. Prog. Solid State Chem. 2004, 32, 33-177.
7. Konstantinou, I. K.; Albanis, T. A. Appl. Catal. B. 2004, 49, 1-14.
8. Hirakawa, T.; Nosaka, Y. Langmuir 2002, 18, 3247-3254
9. Wu, C.; Chang, H.; Chern, J. J. Hazard. Mater. 2006, B137, 336-343
10. Baldini, F.; Giannetti, A. OPTO-Ireland; International Society for Optics and Photonics,
2005; pp. 485–499.
Appendix:
Figure 7. Full wavelength scan between 400 nm and 720 nm of the 5 ppm methylene blue solution. The wavelength
corresponding to the highest absorbance value was observed to be 0.996 at 659 nm.

16
Table 1. Average absorbance and corresponding standard deviation for each methylene blue diluted in increments of
0.1 between 0.1 to 0.9 times the original 5 ppm solution.
Dilution factor of the original 5 ppm
solution
Average
absorbance
Absorbance standard deviation
0.1 0.1140 0.0005
0.2 0.2134 0.0008
0.3 0.3021 0.0007
0.4 0.4117 0.0009
0.5 0.5066 0.0006
0.6 0.6000 0.0009
0.7 0.6974 0.0005
0.8 0.7914 0.0023
0.9 0.8841 0.0007
1 0.9842 0.0027

17
Table 2. Average absorbance and corresponding standard deviation for each sampling time after 5 mL of H2O2 were
added to the TiO2/methylene blue mixture (first trial).
Time
(min)
Average
absorbance
Absorbance
standard deviation
0 0.7184 0.0020
2 0.5093 0.0006
4 0.4762 0.0006
6 0.4511 0.0010
8 0.4818 0.0013
10 0.4421 0.0014
12 0.4133 0.0012
14 0.5966 0.0042
16 0.4427 0.0003
18 0.3451 0.0005
20 0.3307 0.0012
22 0.3342 0.0004
24 0.3337 0.0008
26 0.2949 0.0002
28 0.2683 0.0007
30 0.2592 0.0009
32 0.2492 0.0008
34 0.2377 0.0002
36 0.2114 0.0004
38 0.2035 0.0008
40 0.1905 0.0007
42 0.7072 0.0043
44 0.2102 0.0008
46 0.1559 0.0002
48 0.1526 0.0004
50 0.1358 0.0007
52 0.1269 0.0007
54 0.1221 0.0003
56 0.1034 0.0036
58 0.0906 0.0006
60 0.0835 0.0004
62 0.0725 0.0004
64 0.0710 0.0013
66 0.0538 0.0008
68 0.0394 0.0005

18
Table 3. Average absorbance and corresponding standard deviation for each sampling time after 10 mL of H2O2
were added to the TiO2/methylene blue mixture (second trial).
Time
(min)
Average
absorbance
Absorbance
standard deviation
0 0.6557 0.0022
2 0.4847 0.0010
4 0.3601 0.0003
6 0.3368 0.0005
8 0.3245 0.0007
10 0.3058 0.0013
12 0.2878 0.0012
14 0.2605 0.0006
16 0.2462 0.0007
18 0.2256 0.0008
20 0.2103 0.0008
22 0.1958 0.0008
24 0.1622 0.0007
26 0.1556 0.0004
28 0.1367 0.0005
30 0.1246 0.0007
32 0.1015 0.0006
34 0.0918 0.0004
36 0.0728 0.0005
38 0.0643 0.0005
40 0.0556 0.0003
42 0.0454 0.0002
44 0.0404 0.0006

19
Full derivation of Equation 4 from Equation 3:
= -kC (Equation 3)
1
∫
1
∫
( )
Full derivation of Equation 6 from Equation 3:
= -kC (Equation 3)
( ) = -kC
1 1 1
( )
1
∫
1
∫
1
( )
1
( )
1
( )]|
1
( )
1
(
1
)]|
1
( )
1
(
1
) (
1
)]
1
( )
1
(
1
1
)
1
( )
1
( )
( ) ( )

Final Photocatalysis Lab Report (1) (1)

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