ROBERT A. TAFT WATER RESEARCH CENTER
REPORT NO. TWRC-13
PHOTOLYSIS MECHANISMS
FOR
POLLUTION ABATEMENT
ADVANCED WASTE TREATMENT RESEARCH LABORATORY-XIII
U.S. DEPARTMENT OF THE INTER/OR
FEDERAL WATER POLLUTION CONTROL ADMINISTRATION
OHIO BASIN REGION
Cincinnofi, Ohio
-------
PHOTOLYSIS .MECHANISMS FOR POLLUTION ABATEMENT
by
Layton C. Kinney Victor R. Ivanuski
for
The Advanced Waste Treatment Research Laboratory
Robert A. Taft Water Research Center
This report is submitted in
fulfillment of Contract No.
14-12-433 between the Federal
Water Pollution Control Admin-
istration and the IIT Research
Institute.
U. S. Department of the Interior
Federal Water Pollution Control Administration
Cincinnati, Ohio
October, 1969
-------
FOREWORD
In its assigned function as the Nation's principal natural
resource agency, the United States Department of the Interior
bears a special obligation to ensure that our expendable resources
are conserved, that renewable resources are managed to produce
optimum yields, and that all resources contribute their full
measure to the progress, prosperity, and security of America
now and in the future.
This series of reports has been established to present the
results of intramural and contract research studies carried out
under the guidance of the technical staff of the FV/PCA Robert A.
Taft Water Research Center for the purpose of developing new or
improved wastewater treatment methods. Included is work con-
ducted under cooperative and contractual agreements with Federal,
state, and local agencies/ research institutions, and industrial
organizations. The reports are published essentially as sub-
mitted by the investigators. The ideas and conslusions presented
are, therefore, those of the investigators and not necessarily
those of the FWPCA.
Reports in this series will be distributed as supplies per-
mit. Requests should be sent to the Office of Information, Ohio
Basin Region, Federal Vfeter Pollution Control Administration,
4676 Columbia Parkway, Cincinnati, Ohio 45226.
il
-------
TABLE OF CONTENTS
Foreword ii
Abstract iv
Introduction I
Experimental 3
Screening of Photocatalysts , 3
Photoxidation of Organic Substances 5
Photocatalytic Oxidation of Phenol with TiQ2
and Beach Sand 24
Oxidation of Phenol with the Ace Glass Company
Photochemical Apparatus 29
Oxidation of Raw Sewage 35
Bactericidal Properties of Irradiated Zinc Oxide
Slurries 36
Discussion 38
Future Work 40
Interaction of Ionizing Radiation with Metal
Oxide Catalysts 40
References 41
Appendix Al
111
-------
ABSTRACT
Photocatalytic oxidation of dissolved organic matter by
irradiation of slurries of zinc titanate (Zn2TiO.), zinc oxide
(ZnO), titanium dioxide (TiO2) and beach sand by sunlamps has
been demonstrated. The reaction appears to follow first order
kinetics in most cases. Zinc oxide appears to be superior for
this purpose. At. concentrations of 100 to 200 mg/1 organic
carbon, 80% of phenol, 67% benzoic acid, 44% acetic acid,
40% sodium stearate, and 16% sucrose were oxidized in 24 hrs
with 10 gr/1 zinc oxide catalyst. Continued illumination reduced
organic carbon to a few mg/1 in most cases.
The photocatalytic properties of illuminated beach sand
(87% of phenol removed in 72 hr) strongly suggest that photo-
catalysts are widely distributed in nature and that photo-
catalytic oxidation is a mechanism whereby dissolved organic
matter is oxidized in the natural environment in streams and
lakes..
Dissolved organic matter in a sample of domestic sewage
was reduced 50% in 24 hours and 75% in 70 hours. The limiting
factor appears to be the activity of the photocatalyst since
sunlamp intensity dropped approximately 80% without measurable
decrease in oxidation rate. Improving catalyst activity by
quenching from high temperature and/or by doping is proposed.
The effects of ionizing radiation should be studied
to determine the economic and technical advantages over UV. A
recent publication describing ZnO sensitized polymerization
with both UV and gamma radiation reports that zinc oxide caused
an enormous enhancement of the polymerizing effect of ionizing
radiation in an aqueous system of calcium acrylate and acryla-
mide which suggests a corresponding increase in free radical
generation.
At the present state of knowledge the most promising area
of application appears to be in problems of industrial disposal
where high concentrations of organic matter are involved since
IV
-------
the amount removed at. constant irradiation intensity increases
with increased concentration. Small residual concentrations
might be economically removed by conventional means.
-------
INTRODUCTION
This study was based on photocatalytic phenomena described
in the literature and on experimentation conducted in IITRI
laboratories. The interaction of photocatalysts with radiation
o
below about 4200A should produce active oxygen species which
could destroy organic matter by complete oxidation to carbon
dioxide and water.
Certain oxides, notably zinc oxide and titanium dioxide,
are known to be photosensitizers or photocatalysts. Goodeve and
Kitchener obtained powerful photosensitization with TiO-,
studied by bleaching a blue dye under near-UV irradiation.
2
Jacobsen correlated outdoor chalking behavior of Ti02 pigmented
paints with the loss of reflectance of various TiO2 pigmented
media under sunlamp exposure. rieyl and Forland proposed that
atomic oxygen was dislodged by photons from rutile and other
oxides? e.g., Si02 due to the asymmetric force field acting on
the surface oxygen atoms.
Although photo-oxidation by zinc oxide has been much more
extensively studied than TiO2, no method to make it non-photo-
active for outdoor pigment use has been found. The use of zinc
oxide as a photoconductor in electrophotography has stimulated
4
a great deal of study. Bauer and Neuweiler were first to
report, the formation of hydrogen peroxide in near-UV-irradiated
aqueous suspensions of zinc oxide. Markham et al. studied zinc
oxide photocatalysis extensively in aqueous systems and showed
that sharply increased hydrogen peroxide concentrations were
formed in aqueous media containing organic substances.
Yamamoto and Oster reported that ZnO under UV excitation
causes polymerization of vinyl monomers if water and oxygen are
present. They concluded, on the basis of end group titration
of polymethylmethacrylate polymers formed under these conditions,
that hydroxyl radicals initiate the polymerization.
Work at IITRI showed that UV irradiated aqueous slurries
of a specially prepared TiO2, produce an oxidant similar to the
-------
hydrogen peroxide formed in UV irradiated zinc oxide slurries.
However, this oxidant does not give the amber pertitanic
acid coloration specific for hydrogen peroxide. The formation
of hydrogen peroxide in irradiated zinc oxide slurries has been
confirmed by other investigators by the titanic sulfate reaction.
Both oxidants liberate, of course, iodine from potassium iodide.
It was first thought the unknown oxidant was ozone? however, the
absence of ozone odor at high concentrations and its appreciable
solubility in water (concentrations of 500 ppm have been obtained)
rule out this possibility.
Either oxidant (produced in situ by UV irradiation of
ZnO and/or TiO_) might provide a means of oxidizing dissolved
organic matter in polluted water, since manufactured hydrogen
peroxide and ozone have been used for this purpose.
-------
EXPERIMENTAL
SCREENING OF PHOTOCATALYSTS
Because the level of oxidant obtainable in our previous
work was low (8-10 mg/1), an investigation was undertaken to
determine if greater concentrations of catalyst would produce
higher levels of oxidant.
Initially, the light source was 4 G.E. 100W Fluorescent
Daylight lamps on 4-inch centers in an aluminun foil lined
reflector 18 inches from the sample. Samples of one-liter
volume, in 180 by 100 mm crystallizing dishes, were stirred by
magnetic stirrers separated from the dishes by Coolplates* which
limited heating by the stirrer motors to about 2°C above ambient.
Catalysts initially examined included Photox 801, zinc
oxide. (New Jersey Zinc Co.), titanium dioxide prepared by our own
process, and a sample of zinc titanate. The latter had been
found to be destructive to organic vehicles under UV light.
Initial experiments conducted at 25 ppm verified previous
data and indicated that zinc titanate was a photocatalyst on
the basis of an initial equilibrium oxidant level of approxi-
mately 5 mg/1 as H^O-.
In an effort to raise oxidant concentrations, catalyst
loadings were increased from 1-5%; however, this had the
opposite effect in that equilibrium concentrations dropped sub-
stantially.
We concluded that the catalyst was decomposing the oxidants
and that once the shaded fraction exceeded the irradiated oxide,
the equilibrium concentration dropped.
This effect was quite unexpected, since Markham and Laidler ,
working with fairly concentrated slurries (0.1 g ZnO/25 ml H2O)
had obtained high concentrations (170 mg/1) of H2O2. Their
results are attributable to the fact that the various organic
substances they worked with inhibited the decomposition reaction,
*Thermodyne Corporation Dubuque, Iowa.
-------
thus strongly favoring oxidant formation.
Since the dilute slurries absorb only a small fraction of
the radiation, use of an oxidized metal surface as the photo-
catalyst was investigated. In this way access of the solution
to shaded surfaces would be decreased or prevented. Also, in
many cases, it was thought the oxidized surface would be a more
efficient absorber of radiation than the white pigments. On
this assumption a number of metals were examined including zinc,
copper, nickel, bronze, titanium, and tin.
In some cases the metal was heavily oxidized by heating
and in others it was used with the normal invisible oxide of
the metal. Many of these materials, especially zinc and copper,
looked quite promising at first, on the basis of oxidant levels
which in the case of zinc rose rapidly at first to about
25 mg/1 H2°2' nowever' as the formation of zinc oxide proceeded
it dropped to 2.5 mg/1 in about 3 days.
In addition to oxidized metals it was found that a number
of other substances show photocatalytic properties, for example,
Ottawa silica sand and two man-made pigments which were largely
Fe2O3 and Fe3O4. Two dimensional catalyst surfaces such as
ceramic porcelain ware and Pyrex as well as films of zinc
undecylate developed oxidant concentrations of f-10 mg/1
when illuminated for a week or so. Of the oxidants formed, all
were of the non-hydrogen peroxide variety (with the exceptions
of irradiated zinc and tin metal plates) as determined by their
negative pertitanic acid tests with titanic sulfate reagent.
Since we have not yet identified oxidant "x", as we call it,
and since it is detected by the oxidation of potassium iodide
these oxidants were both reported as hydrogen peroxide as deter-
mined by a spectrophotometrie starch iodide method described in
detail in the Appendix.
»fe concluded that there must be many substances, other
than man-made, which exhibit photocatalytic properties; in all
probability ordinary beach sand contains photocatalytic minerals.
To test this hypothesis, sand was obtained from a nearby beach
-------
and exposed to daylight fluorescent radiation.
After 2 days of exposure, of a thin layer of sand, the oxidant
level rose to 3.5 mg/1. A thick layer of sand produced a faint
but definite trace of oxidant which remained constant after a
few days of exposure. The difference is attributed to the
enhancement of the reverse reaction in the thick layer. Appar-
ently, our hypothesis was correct. Oxidant "x" is being formed
and decomposed on shores and beds of streams where natural
photocatalysts are illuminated by solar energy. We theorized
that photocatalysis is a mechanism of nature to remove organic
pollutants in addition to non-catalytic photolysis, autoxidation
and biodegradation. This theory was subsequently verified by
brief, but conclusive, experiments demonstrating the removal of
organic contaminants from water by simulated sunlight. These
experiments are described in the following section.
The site from which sand samples were obtained was sub-
stantially that from which PettiJohn obtained his. Since this
is natural beach, it is assumed the mineral content is similar
to that he reports. The principal minerals found by PettiJohn
were garnet (Mg, Ca, or Fe-Al silicate); augite (CaO'MgO-2SiO2);
Hornblende (3lMgFe]O'CaO); Si02/- ilmenite (FeO 48%, TiO2 52%), and
magnetite (Fe^O.) along with calcium and magnesium carbonaceous
minerals.
PHOTOXIDATION OF ORGANIC SUBSTANCES
Although we had hoped to find a catalyst superior to the
powdered oxides initially investigated, it became increasingly
evident that oxidant level, per se, was not a valid indication
of organic oxidation effectiveness. Consequently, efforts were
directed toward actual oxidation studies which were monitored
by the Beckman Total Organic Carbon Analyzer.
The first experiments/ run in zinc pans, were inconclusive.
Although linear alkylate sulfonate (LAS) was oxidized, we sus-
pected that galvanic factors due to the formation of oxygen
concentration cells were involved since oxidation in dark samples
-------
approached that of illuminated samples. For this reason further
work with oxidized metal surfaces was abandoned.
The first definite photo-oxidation of an organic substance
obtained is shown in Figure 1. The apparent initial rise in
organic carbon is not understood; it may have been due to
desorption of an organic substance from the catalyst (20 g zinc
titanate in 350 ml H9O). The superiority of the sunlamps* whose
^ o
emission is almost entirely in the 3000-4000A region over the
G.E. Daylight fluorescent lamps, is clearly evident from the
Figure* This superiority was somewhat surprising since oxidant
levels were about the same under each source. This demonstrates
that oxidant level alone is not a valid indication of organic
oxidation rate. With few exceptions, subsequent tests were run
under sunlamps. These samples were exposed in 180 mm-diameter
Pyrex dishes, 75 mm deep without agitation. In view of its poor
showing in these tests, further work with ZnTiO. was discontinued.
Figure 2 demonstrates oxidation of both phenol and benzoic
acid and the H0O0 concentration does increase in the presence of
5
organic matter as reported by Markham and Laidler . The higher
values obtained by those authors are probably due to the fact
they used pure oxygen. Markham and Laidler, as well as Oster
Q
and Yamamoto , showed oxygen to be essential to the photo-oxida-
tion by zinc oxide.
The samples in Figure 2 were exposed in 180 x 100 mm
crystallizing dishes and stirred magnetically. Dishes were
placed eccentrically which helps to prevent the catalyst from
settling. Analyses for total organic carbon (TOC) and oxidant
were made on "whole samples", i.e., catalyst and solution which
avoids artifacts due to adsorption of degradation products on
the catalyst which appears to vary in the course of the test.
In hundreds of tests, the small quantity of oxide introduced
into the combustion zone of the organic carbon analyzer had no
apparent effect.
___ __ ___
8 twenty watt sunlamps on 4" centers with an aluminum reflector
18" from the sample.
-------
10,000 r
1000
A 1 x G
O 1 x a
D i x o
2 x G
2 x S
V 2x0
x = 100 mg/1 of LAS.
G indicates daylight fluorescent (G.E.)
S indicates sunlamp (Westinghouse)-
D indicates dark.
I
10
345678 9 !0
Time, days
Figure 1
EFFECT OF LIGHT SOURCE ON LAS REMOVAL
BY ZINC TITANATE
11
12
13
-------
00
Table 1
EFFECT OF LIGHT SOURCE ON LAS REMOVAL BY ZINC TITANATE
Concentration of Organic Carbon, mg/1
Sample
IxG
IxS
IxD
2x3
2xS
2xD
Tlours .
0
Carbon
35.0
35.0
35.0
69.5
69.5
69.5
>
J9H-
7.2
7.2
7.2
7.0
7.0
7.0
t = 100 n
3a
Carbon
47.5
48.5
45.0
76.0
74.0
74.0
ng/1 of
Time, day
1 2
pH Carbon pH Carbon pH
6.1 47.5 48.5
6.3 50.0 51.0
6.4 50.0
6.2 81.0 80.0
6.2 80.5 81.0
6.4 66.0
LAS.
9 11
Carbon pH Carbon pH
36.0 5.9 36.0 6.0
26.0 6.1 14.5 6.0
39.0 5.8
69.5 4.8 65.0 5.4
53.5 4.7 28.0 5.8
62.0 5.8
G indicates daylight fluorescent (G.E.)
S indicates sunlanp (Westinghouse).
D indicates dark.
-------
10,000
10,000
Symbol
1,000
H2°2
Organic
Carbon
o
A
D
O
Sample
Designation
1*3 PL
3 PL
3 PD
BL
Contaminant
Phenol (light)
Phenol (light)
Phenol (dark)
(control)
Benzoic acid
(light)
Tapwater
1 liter
l*j-in. deep
2 liter
3-in. deep
2 liter
3-in. deep
1 liter
1*1-in. deep
1,000
Note: Catalyst, 5 g/1 Photox 801, ZnO; illumination,
sun lamps
u
o
3 PD
100
10
10
20 30
40 50
100 110 120 130
60 70 80
Time, hr.
Figure 2
EFFECT OF DEPTH OF SAMPLE ON OXIDATION OF PHENOL
AND COMPARISON OF PHENOL AND BENZOIC ACID OXIDATIONS
£
01
O
n
K
m
O
c
0
-H
.p
a
£
u
8
-------
Table 2
EFFECT OF DEPTH OF SAMPLE ON OXIDATION OF PHENOL
AND COMPARISON OF PHENOL AND BENZOIC ACID OXIDATIONS
(Catalyst, 5 g/1, Photox 801 (N.J.Z.)
Concentration of Organic Carbon and H0O,, mg/1
Time, hr
0
3
18
19
20
22
25
41
44
45
68
92
112
Sample
l-l/2aPL 3 PL 3 PDb
Organic Organic Organic
Carbon M2°2 Carbon H2°2 Carbon H2°2
122 104 114
106
106
73 38
15 35
104
44
11 3.5
14
6 23
7 8
1-1/2 BL
Organic
Carbon M2°2
101
27 56
9 32
3 16
1-1/2-in. depth, 3-in. depth, etc.
Control sample kept in dark.
Note: The phenol sample concentrations were approximately 0.15 g/1,
and the benzole acid sample concentration was also approximately
0.15 g/1.
Sample
Designation
lh PL
3 PL
3 PD
BL
Contaminant
Phenol (light)
Phenol (light)
Phenol (dark)
(control)
Benzoic acid
(light)
Tapwater
1 liter
1*5-in. deep
2 liter
3-in. deep
2 liter
3-in. deep
1 liter
1*5-in. deep
10
-------
First order kinetics were observed for the oxidation of
both phenol and benzoic acid. The decrease in hydrogen peroxide
concentration after a sharp initial rise appears to follow a
similar path, though the data are incomplete. Control samples
stored in the dark showed no reaction between either photo-
catalytically produced oxidant or H2O_ with organic matter.
This suggests that the oxidation is directly associated with the
active oxygen species released by the radiation.
Figure 3 represents an effort to assess the importance of
catalyst loading on phenol oxidation. Samples were magnetically
stirred in 100 mm by 180 mm crystallizing dishes. The data
indicate only slight gains by increased catalyst loading from
2.5 g/1 - 10 g/1, whereas at 0.5 and 0.1 g/1 proportionately
lower oxidation rates were encountered. This suggests that once
all the incident light is absorbed, additional catalyst gives
only slightly improved removal of organic carbon.
Despite the apparent lack of effect of excess catalyst,
a recirculation system was devised to keep all the catalyst in
suspension.
Several three gallon Pyrex bottles were cut off 2 inches
from the mold line at the shoulder and the bottom was discarded.
The neck was reduced in diameter and a 29/42 standard taper
female fitting was fused to the reduced neck. A male taper
fitting with a suitable opening for drainage was attached by
springs and the assembly was inverted on an iron tripod. The
bottom drain was connected to a the suction side to a peris-
taltic pump with a Viton tube after first passing the slurry
through a water-cooled condenser. The discharge of the pump
was connected to a 10 mm tube with a right angle bend terminating
in a 3 mm nozzle which was arranged at the outer periphery of
the vessel so that the discharge imparted a circular motion to
the contained slurry. In this manner the mixture was constantly
recirculated at a rate of about 1 1/min. Additional aeration
was provided by a sintered glass gas dispersion tube immersed to
within about 1 inch from the pump intake. Air from an aquarium
11
-------
1000
1000
0.1P
* 100
I _
3 h-
-------
Table 3
EFFECT OF VARIOUS ZINC OXIDE LEVELS ON MAGNETICALLY STIRRED SAMPLES
Sample
10P
5P
2.5P
0.5P
0.1P
(cont.)
10P
5P
2.5P
0.5P
0.1P
Time,
hr
0
0
0
0
0
21.7
24
21.5
21.5
21.7
Concentration,
ma/1
Organic H n
Carbon 22
161
150
155
161
155
68 48
83
71
72
111 30
Concentration,
mg/1
Time, Organic H n
hr Carbon 2U2
2.3 152
4 150
2.2 162 22
2.2 154
2.2 152 15
24 61 50
51a
29 63
29 82
29 103 36
aCentr if uged .
Concentration,
ma/1
Time, Organic H n
hr Carbon T2
4 29.6
3.5 28.5
30 45 38
Sample
Designation
10P
5P
2.5P
0.5P
0.1P
Concentration, Concentration,
ma/1 ma/1
Time, Organic n Time, Organic H o
hr Carbon 22 hr Carbon 2y2
5.8 144 35 10.7 118
7 137
5.4 151 10.7 125
5.6 149 10.7 118 28
5.8 153 22 10.5 134
48 16
6a
48 19
48 23
Concentration
of znO, g/la
10
5
2.5
0.5
0.1
-------
pump was provided at a rate of 60 ml/min.
Figure 4 depicts results of recirculation at various cata-
lyst loadings and shows a sharp increase at 10 g/1. At catalyst
loadings of 2 and 5 g/1, however, there seems little effect.
Even with the recirulcation system sedimentation of the catalyst
occurred in certain instances. It is interesting to note the
behavior of the H2O_ concentration in these tests. With phenol
contaminant the level drops more rapidly with higher catalyst
loading and shows a general relationship to the residual contam-
inant. This does not hold true for acetic acid however, since
the concentration drops sharply while the residual contaminant
is still high. The H2O2 concentration for sucrose (Figure 7)
appears to level off and remain high despite the high catalyst
loading. A better understanding of these phenomena might provide
useful insights into the effects of pollutants in the natural
environment.
Figure 5 shows little difference between magnetic stirring
or recirculation if one takes into account the lower starting
concentration of samples 5P, iJjBl, and lljPl. The curve for the
1%P1 sample may be the result of sampling errors.
Figure 6, on the other hand, shows a clearcut advantage
of recirculation over magnetic stirring. Figure 7 shows a com-
parison of phenol with sucrose, sodium stearate and two examples
of acetic acid oxidation. The difference in the two acetic acid
runs is due to the sequence of adding the zinc oxide and acetic
acid. The acetic acid curve on Figure 4 was obtained by first
adding the acid to tap water and then adding the zinc oxide.
This method removes all of the inorganic carbon. However, over
the course of the oxidation inorganic carbon returned to the
original value. In a later experiment, shown in Figure 7, the
zinc oxide was added before the acetic acid, which results in
the formation of zinc acetate which is more slowly oxidized,
accompanied by a substantial rise in inorganic carbon means
that one product of oxidation of zinc acetate is zinc carbonate.
From Figure 7, it can be seen that phenol, acetic acid (added
14
-------
1000
Symbol
0 organic
22 Carbon
i i
U
- A A
O
+ 0
Sample
Desiqnat.
10 AP
5 AP
2 AP
1 AP
Concentration
Lon Contaminant of ZnO, g/la
T5Vii=ini-»l 10
Phenol 5
Phenol 2
Phenol 1
V ^7 10 HAC Acetic Acid 10
\,iter
tapwater
1000
Note:
Illumination, 8 20-w Westinghouse
sunlamps at 18-in. on 4-in. centers
of aluminum foil reflector
100
00
I
3
U
M-l
O
O
H
§
O
c
8
12 18 24 30 36
Figure 4
42
54
60
EFFECT OF ZINC OXIDE LOADING
ON RECIRCULATED SAMPLES
15
-------
Table 4
EFFECT OF ZINC OXIDE LOADING ON RECIRCUIATED SAMPLES
Sample
10 AP
r AP
2 AP
1 AP
10 HAC
( cont . )
10 AP
5 AP
2 AP
10 HAC
Time,
hr
0
0
0
0
24
29.7
28.8
30
Concentration.
mg/1
Organic o
Carbon "22
157.5 3.5
190
160 3.5
170 3.5
178
20 4«
60 60
70 60
95 42
Discontinued.
Concentration, Concentration, Concentration, Concentration,
mg/1 mg/1 mg/1 mg/1
Time, Organic
hr Carbon "
.75
1 183
.3
2 139
3 156
27 14
4ft 27
4fc 10
46 bS
Sample
Designation
TOAP
5 AP
2 AP
1 AP
10 HAC
n Time, Organic u ... . Time, Organic _ Time, Organic u _
2U2 hr Carbon H2°2 hr Carbon H2°2 hr Carbon - H2°2
21 1.5 141 4 126 35 9.5 82 40
3.1 170 31 5.3 146 37 22 82 66
12 2.4 175 5.2 147 38 22 88 57
30 5.3 115 39 18.7 48 42 24 43 52*
30 5.5 148 37 23 124 60
50 30 10 53 49 7 36
47
54
50
Concentration
ortaminant of ZnO, g/la
:jhenol 10
Flienoi. 5
Phenol 2
Phenol 1
Acetic Acid 10
-------
1000
100
u
o
H
s
o>
o
4J
0)
O
O
O
Sample
Symbol Designation Sample Reference
D 5P Phenol Table 3
A
1*5 PL
BL
Phenoj
Benzoic
Acid
Phenol
Table 2
Table 2
Table 4
Method of
Agitation
Magnetic
Stirring
Magnetic
Stirring
Magnetic
Stirring
Recirculation
10
Note: Tapwater used for ZnO solution; illumination, 8 20-w
Westinghouse sunlamps at 18-in. on 4-in. centers of
aluminum foil reflector.
12 18 24 30 36 42
Time,hr.
Figure 5
COMPARISON OF AGITATION METHODS
USING A ZnO CONCENTRATION OF 5 g/1
48
54
60
17
-------
1000
Symbol
a
o
Sample
Designation
10 A*
10 P
Reference
Figure 4
Figure 3
Method of
Agitation
Recirculation
Magnetic
Stirring
Note: Illumination, 8 20-w Westhinghouse sunlamps
at 18-in. on 4-in. centers of aluminum
foil reflection
Figure 6
COMPARISON OF AGITATION METHODS ON OXIDATION OF PHENOL
USING 10 g OF ZINC OXIDE
18
-------
1000
000
Symbol
H2°2
Organic Sample
Carbon Designation
D
A
O
d
0
A
10 s
5 P
10 HAC
5 AP
10 HAC
10 SS
Sample
Sucrose
Phenol
Refer- Concen.
ence ZnO, g/1
Treatment
Acetic
Acid
Phenol
Acetic
Acid
Sodium
Stearate
Fig.
Pig.
10
5
10
10
10
Recirculation
Recirculation
Recirculation
RecirculationJ
Aeration
Recirculation_
Aeration ~
Recirculation
D
10 S
Note:
ZnO solutions made with tapwater, except for
10 SS sample for which distilled water was used;
illumination, sunlamps.
00
-------
Table 5
(Data plotted in Figure 7)
COMPARISON OF AERATED WITH NONAERATED RECIRCULATED SLURRIES,
ALSO OXIDATION OF ACETIC ACID SUCROSE, PHENOL, AND SODIUM STEARATE
Sample
10Sa
5Pa
10HAC8
iossa
(cont.)
10S
5P
10HAC
Concentration
Contaminant of ZnO, a/1
Sucrose 10
Phenol 5
Acetic 10
Acid
Sodium 10
Stearate
Concentration
ma/1
Time, Organic n
hr Carbon n2°2
0
0
0
0
27
27
296
152
166
69
.7 249 55
.3 80
28.5 118
Concentration
ma/1
Time, Organic _
hr Carbon 2°2
3.5 276 25
4.3 139
5.0 166
4.1 67 8
50 228 0 2
50 18
51 97
Time,
hr
5.3
__
25
70
70
70
Concentrat ion
ma/1
Organic H Q
Carbon H2°2
289
__
42
188
Concentration
Time, Organic
hr Carbon
22.9 252
22.3 92
22.7 130
discontinued
46 133
48
60
discontinued
89
96 64
aSamples recirculated without aeration.
Samj
Desiai
10
5
10
pie Refer- Concen.
nation Sample ence ZnO, a/1
S Sucrose 10
P Phenol 5
HAC Acetic 10
Acid
Treatment
Recirculation
Recirculation
Recirculation
Sample
Desiqnation
5 AP
10 HAC
10 SS
Sample
Phenol
Acetic
Acid
Sodium
Refer- Concen.
ence ZnO, g/1
Fig. 4 5
Fig. 4 10
10
Treatment
Recirculation
Aeration
Recirculation
Aeration
Recirculation
Stearate
-------
before zinc oxide), sodium stearate and sucrose are oxidized with
increasing difficulty in that order. The sodium stearate curve
represents zinc stearate oxidation. Figure 7 also compares re-
circulated aerated samples with non-aerated recirculated samples.
These results demonstrate that aeration does not enhance
the oxidation mechanism. No differences in dissolved oxygen
were noted, being in both cases, saturated at the temperature of
tv.e tests.
To obtain a more coherent picture of the effects of agita-
tion as well as to assess the degree to which air stripping and/
or autoxidation are involved, the experiments illustrated in
Figure 8 were conducted.
This figure shows a comparison of various methods of agita-
tion with respect to phenol removal. Comparison of the aerated
samples in the dark and the illuminated samples, without catalyst
in each case, shows that part of the phenol removal is by strip-
ping and/or autoxidation. The additional removal under illumin-
ation (without ZnO catalyst) are photocatalytic oxidation by
Pyrex which is a weak photocatalyst*. A similar acceleration
occurs with magnetically stirred samples in the light without
catalyst compared to similar samples in the dark. Phenol added
to distilled water stirred in polyethylene (not shown) had a
slope closely paralleling the magnetically stirred sample in
the dark.
Since much of the increase in oxidation rate was due to air
stripping in both the recirculated and the aerated samples, we
decided to return to magnetic stirring because this method shows
the greatest effect due to the presence of catalyst. However,
none of the methods keeps the catalyst entirely in suspension.
Initially we thought as long as sufficient catalyst was in sus-
pension to absorb all the incident radiation, an excess would do
no harm. This is true if the pigment all remains in suspension;
*Distilled water magnetically stirred reaches an equilibrium
concentration of oxidant "x" as I^O^of 4 mg/1 when illuminated
by G.E. Daylight fluorescent lamps for 6 days.
21
-------
1000
Sample
! Symbol Designation
10 AP
10 P
APL
Comment
o
- o
A
RPL
APD
MPL
1-1PD
10 g of znO; recirculated; from
Figure 6.
10 g of ZnO; magnetically stirred; from
Figure 6.
Aerated 4 liter/min through 1-mm
capillary in bottom of inverted
3-gal bottle; no catalyst; sunlamp
illumination.
Same container; recirculated
1 liter/min tangential discharge; no
catalyst, sunlamp illumination.
Same as APL, but no illumination.
Magnetically stirred, 180-mm x
100-mm Pyrex crystallizing dish;
no catalyst; sunlamp illumination.
Same as 1. '_, but no illumination.
100
s
o>
Ll
-U
IB
O
8
10
18 60
Time, hr.
108
120
Figure 8
COMPARISON OF VARIOUS AGITATION METHODS
IN REMOVING PHENOL
22
-------
10
U)
Table 6
(Plotted in Figure 8)
COMPARISON OF METHODS OF AGITATION ON OXIDATION OF PHENOL
USING lOg OF ZINC OXIDE
Concentration Concentration
ma/1 ma/1
Time
Sample hr
APL 0
RPL 0
APD 0
MPL 0
MPD 0
Sample
Designation
APL
, Organic
Carbon
154
153
140
155
157
n Time, Organic n Time,
"2U2 hr Carbon H2°2 hr
23 112
23 101
24 97
19 152
19 155
Comment
Aerated 4 liter/min through 1-mtn
capillary in bottom of inverted
3-gal bottle; no catalyst; sunlamp
48
48
48
1.5 43
43
Sample
Designation
APD
MPL
Concent rat ion
ma/1
Organic n Time
Carbon 22 hr
52 72
67 3.0 72
70 72
115 96
140 96
Comment
Same as APL, but no
Magnetically stirred
Concentration
ma/1
, Organic n
Carbon 22
50
51
85
112
illumination.
, 180-mm x
RPL
illumination.
Same container; recirculated
1 liter/min tangential discharge;
no catalyst, sunlamp illumination.
MPD
100-mm Pyrex crystallizing dish;
no catalyst; sunlamp illumination.
Same as MPL, but no illumination.
-------
however, if the contaminant is appreciably adsorbed on settled
catalyst, it is protected from oxidation and may only become
available for oxidation through diffusion, thereby slowing the
process.
PHOTOCATALYTIC OXIDATION OF PHENOL WITH Ti02 AND BEACH SAND
Since we had on hand only limited amounts of the special
TiO~ catalyst prepared in our previous work, it was necessary to
prepare it. The Ti02 was obtained from New Jersey zinc Co. in
slurry form and did not contain the stabilizing additives, e.g.,
silicates that are normally added to minimize photocatalysis in
commercial pigments. This slurry was washed chloride-free and
nearly sulfate-free. The resultant cake was dried at 100°C.
Figure 9 compares oxidation of phenol by various types of TiO2.
Based on limited tests the optimum calcination temperature is
about 850 °C since at 48 and 72 hrs, the 850°c sample is better
than either the 825°C or 875°C sample. The apparent increase in
rate of oxidation with time is interesting. In subsequent tests
(not shown), a commercial anatase and an active TiO2 prepared
some years ago by the IITRI technique were also compared under
the same conditions; both were definitely inferior to the samples
shown in Figure 9 in oxidation of phenol under the same conditions
A brief experiment was conducted to determine the effec-
tiveness of beach sand. These data are shown graphically in
Figure 10 and support our previously stated conviction that
photo-oxidation by natural catalysts under solar irradiation is
a mechanism of nature for the removal of organic matter from
streams and lakes. In these tests, 300 g of water washed beach
sand were placed in the recirculation container described pre-
viously and 4 1/min of room air was pumped through from the bot-
tom. This produced much less agitation in sand than anticipated
and much less than sample APD, since the column of water being
agitated was about six times longer due to the absence of sand.
The true comparison is between 300 PBD and 300 PEL. The leveling
off of 300 PEL at 22 mg/1 is believed due to the presence of a
refractory organic substance leaching out of the sand or possibly
24
-------
1000
100
I
id
U
n)
o
.u
-------
Table 7
(Plotted in Figure 9)
OXIDATION OP PHENOL WITH ANATASE AND ANATASE-RUTILE
PREPARED BY CALCINATION AT DIFFERENT TEMPERATURES
Sample
1-AP
1-825 ARP
1-850 ARP
1-875 ARP
( cont . )
1-825 ARP
1-850 ARP
1-875
Catalyst
Anatase
Anatase-
Rutile
Anatase-
Rutile
Anatase-
Rutile
Calcination Time,
Temperature °C hr
Not calcined 0
825 0
850 0
875 0
72
72
72
Concentration
ma/1
Organic Q
Carbon 22
155 0
166 0
168 0
156 0
39
32
49
Concentration
ma/1
Time, Organic _
hr Carbon n2u2
25.7 120 0
26 110 3.3
26 105 3.5
23 137 2.8
96 24
Concentration
ma/1
Time, Organic
hr Carbon
48 89
48 70.0
48 67
50 75
H2°2
3.0*
Oxidant, reported as H000
fc £*
Note: All samples contained 1 g of pigment
-------
1000
o
300 PBD
100
o
o
H
10
D>
O
rt
4J
a
M
4J
V
O
§
10
Sample
Symbol Designation
300 PB
Sample
D
O
o
A
300 PBD
APD
PEL
Aerated
Beach Sand
Aerated
Beach Sand
Aerated
Beach Sand
(no catalyst)
control
Polyethylene
dish, uiagnet-
ically stirred
(no catalyst)
Illumination
Sunlamp
None
None
Sunlamp
Note: Contaminant, phenol.
12
24 36 48 60 72 84 96
Time, hr.
Figure 10
EFFECT OF ILLUMINATION ON OXIDATION OF PHENOL
BY BEACH SAND
108
120
27
-------
to
00
Table 8
(Plotted in Figure 10)
EFFECT OF ILLUMINATION ON OXIDATION OF PHENOL BY BEACH SAND
r
Sample
300 PEL
300 PBD
PEL
( cont . )
300 PEL
300 PBD
PEL
Sample
Designation
300 PEL
300 PBD
Concentration
ma/1
rime, Organic n Time,
hr Carbon n2 2 hr
0 172 3.1
0 185 0 16
0 181 22
42.5 54 72
72 134 96
72 96
Sample Illumination
Aerated Sunlamp
Beach Sand
Aerated none
Beach Sand
Concentration Concentration
ma/1 mq/1
Organic _ Time, Organic
Carbon 22 hr Carbon
187 24 132
144 . 48 131
149 48 124
23 11* 94 23
96
Sample
Designation Sample
APD Aerated
Beach Sand
(no catalyst)
control
H202
13*
Illumination
none
PEL
Oxidant, reported as HO
Polyethylene
dish, magnet-
ically stirred
(no catalyst)
Sunlamp
-------
to particulate organic matter. All samples were whole samples,
i.e., they included any participates generated by the agitation.
Sample PEL shows that the oxidation is primarily photocatalytic
and not simply due to photolysis.
o
The intensity of 3000-3800A of sunlight is somewhat lower
than that of the sunlamps so the natural process may be somewhat
slower. On the other hand, in the range of about 3800-4200A,
sunlight is much stronger than the sunlamps, thus the oxidation
rates of sunlamp vs. solar irradiation is a question of the
relative effectiveness of radiation bands. Unfortunately time
did not permit further exploration of this matter.
OXIDATION OF PHENOL WITH THE ACE GLASS COMPANY PHOTOCHEMICAL
APPARATUS
The photochemical apparatus consists of a medium-pressure
mercury arc contained in a water-cooled quartz jacket which in
turn is immersed in a Pyrex jacket. The sample to be irradiated
is introduced into the annular space between the quartz and
Pyrex containers, where it receives virtually all of the irra-
diation. Provision is made for sparging, and the exposure space
may be pressurized.
o
The 450-watt mercury arc lamp provides about 0.063 watts/cm
2
of effective illumination vs. about 0.0032 watts/cm for the
sunlamps. The increase in oxidation rate due to the higher
intensity was quite modest, as can be seen from comparing Fig-
ures 8 and 11. Figure 11 also represents an attempt at acceler-
ation of oxidation by dye sensitization which is widely used to
broaden the range of spectral absorption in electrophotography.
Without a Pyrex filter (sample RB#1 Lrose bengalj)(Figure 11)
the dye is bleached in a few minutes. We believe RB#2 (Figure 11.)
represents true dye sensitization but the dye oxidized too
rapidly for dye sensitization to be of practical interest.
Sudan orange is much more oxidation-resistant than rose
bengal but does not increase photosensitivity. The use of high
29
-------
1000
Sample
- Symbol Designation Sample Catalyst (60q) Sensitizer Temp.,°c
Cl
|
u
u
H
10
01
A
D
O
O
Phenol Photox 801
(control)
None
40(aver.)
RB-rfl
RB42
R1V3
3041
Phenol
Phenol
Phenol
Phenol
Photox 801 Rose Bengal 22
dye
Photox 801 Rose Bengal3 22
dye
Photox 80 Rose Bengal* 22
dye
Photox 801 Sudan Orangea 22
dye
aPyrex filter.
100 _
X-l
O
01
O
H
4J
(0
k.
P
IB
U
c
O
10 _
RB42
10
20
30 40
Time, hr.
50
60
70
Figure 11
EFFECT OF DYE SENSITIZER ON PHOTOCATAJLYTIC OXIDATION
IN 'ACE GLASS PHOTOCHEMICAL APPARATUS
30
-------
Table 9
(Data plotted in Figure 11)
EFFECT OF DYE SENSITIZER ON PHOTOCATALYTIC OXIDATION
IN ACE GLASS PHOTOCHEMICAL APPARATUS
Sample
Cl
RB#1
RB#2
RB#3
S0#l
(cont.)
Cl
RB#2
RB#3
S0#l
Concentration
mtj/1
Time, Organic n
hr Carbon 22
0 213
0 184
0 184
0 215
0 248
11 46
8.5 41 58
24 19
22.5 15
Sample
Designation sample Catalyst
Cl Phenol Photox
(control)
RB#1 Phenol Photox
Concentration,
ma/1
Time, Organic n Time,
hr Carbon 22 hr
0.5 219 1.0
0.25 180 25 3
0.5 184 3
0.16 189 1.3
0.75 227 2.6
.13. ::> 19 16.5
22. > 10
(60g) Sensitizer Temp., °C
801 None 40 (Aver.)
801 Rose Bengal 22
Concentration ,
ma/1
Organic _ Time,
Carbon 22 hr
195 2.75
130 24
133 56 4.5
169 46 3.4
201 6.1
2
Sample
Designation Sample
RB-f*2 Phenol
RB#3 Phenol
SO#1 Phenol
Concentration,
ma/1
Organic H Q
Carbon 22
177
14 44 (8
88 62
135 56
118
Catalyst (60g)
Photox 801
Photox 80
Photox 801
Concentration ,
ma/1
Time, Organic H Q
hr Carbon 2 2
6.5 122
centrifuged)
5.5 58 60
5.4 110
9.75 93
Sensitizer Temp. , °C
Rose Bengal3 22
dye
Rose Bengala 22
dye
Sudan Orange3 22
dye
-------
loadings of ZnO (60 g/1) was an attempt to compensate for the
short path exposed to the radiation, approximately ^ inch in the
photochemical apparatus as compared to 1^ inch for the magneti-
cally stirred samples. This may have been ill-advised since
data at 10 and 20 g ZnO per liter showed somewhat faster oxida-
tion rates. However, it was assumed that since 20 g/1 was better
than 10 g/1 (not shown) still higher loadings would be better.
The 10 g/1 ZnO sample dropped to 10 mg/1 organic carbon in 12
hours whereas the 20 g ZnO/1 dropped to 10 mg/1 organic carbon
in 8 hours. Comparable time for 60 g/1 ZnO loading from Fig-
ure 12, sample Cl, 10 mg/1 organic carbon was reached in 14 hours.
All tests started at about 220 mg/1 organic carbon. The temper-
ature of the solution in these runs fluctuated from 30° to 50°C.
RB#3, on Figure 11, is a comparison of Photox 80, a zinc
oxide reputed to be about twice as sensitive as Photox 801 in
the visible when dye sensitized and is directly comparable to
RB#2. It appears that this oxide is perhaps less effective
than Photox 801 under these conditions. From this it is evident
there is only a general relationship (if any) between the photo-
conductor properties of the pigment and its photocatalytic prop-
erties, Thus, Photox 801 may have been a poor choice. In view
of the very large number of variables, it was imperative to
standardize on one ZnO catalyst.
After brief raw sewage oxidation studies, described sub-
sequently, work with the Ace glass equipment was discontinued,
It was only recently discovered that the catalyst had a tendency
to adhere to the quartz envelope during the course of the oxida-
tion. While this did not appear to block much light, the adsorp-
tion of phenol oxidation products caused a black tar to form
in the early stages which retards the oxidation to an unknown
degree.
The tar film was discovered when a test was halted in the
early stages of oxidation. It is not known if this was a factor
in the sewage oxidation studies.
32
-------
100
u>
u>
n)
U
o
H
C
c
o
H
-p
(0
M
P
c
0)
o
c
o
U
Benzole acia added (22 mg/1 of organic carton)
Ferroas a;u\.or\i 'rv svilfr-te hexa^ydrate added
'0.162 no)
Sy.nibql Sample
o
Note:
Organic
Carbon
H2°2
Illumination, Westinghouse"
sunlamps; samples magneti-
cally stirred.
100
10
40
60 ~ 80
Time, hr.
Figure 12
OXIDATION OF RAW SEWAGE
100
120
140
CNJ
O
C
n:
o
-H
V
m
c
0)
o
c
o
U
-------
Table 10
OXIDATION OF RAW SEWAGE
(Mixed industrial and
domestic, Park Forest, 111.)
Time,
hr
0
After catalyst
addition
24
24
After benzoic
acid addition
28
27
45
50*
51
56
70
Concentration,
Organic
Carbon
80
28a
28a
40
60
63
60
42
37
31
20
ma/1
H2°2
;
3.5
9.0
22
13
0
0
0
not reliable &.s explained in the text.
bAdded Fe2+ catalyst
34
-------
OXIDATION OF RAW SEWAGE
One experiment,(Figure 12), was conducted with mixed sani-
tary and industrial sewage magnetically stirred under sunlamps.
Although our initial organic carbon oxidation data at 5 minutes
and 1 hour are now known to be faulty, the experiment did repre-
9
sent an attempt at homogeneous catalysis. Bishop, et al. have
shown that ferric and ferrous ion catalyze H2O2 decornP°sit:Lon
with formation of hydroxyl radicals, which, at about pH 4, oxi-
dizes substantial amounts of organic carbon.
Chambron and Giraud reported up to 98% catalytic oxida-
tion of sodium lauryl sulfate by H2O_ with cupric, ferrous and
manganous ion. However, they used H202 concentrations of
approximately 3 wt %.
In our experiment, mixed industrial and domestic raw
sewage was filtered through S & S 597 filter paper and the
organic carbon was determined at the outset. This value started
at 80 mg/1 organic carbon and dropped almost immediately to
28 mg/1 organic carbon. This rapid drop was subsequently proved
to be a sampling error. This sewage contained varying amounts
of "soluble" oil as evidenced by its typical opalescent appear-
ance. In tests of known "soluble" oil concentration, ZnO breaks
the emulsion and allows the oil to float on the surface. This
fact is obscured by the zinc oxide; unless great care is taken to
agitate the sample thoroughly, wide variations occur. In the
case of the actual sewage, a fortuitous check value at 2 hours
gave the reading undeserved credency, initially. For this
reason, we have omitted data prior to the 22 hour point, which
we believe to be valid, as well as subsequent measurements,
since the oxidation follows the usual first order kinetics. The
principal objective here was an attempt to produce hydroxyl
2+
radical by decomposing H2O2 with Fe catalyst after first
enhancing HJX formation by the addition of benzoic acid.
Although, as can be seen from Figure 12, the H,O, dropped imme-
2+
diately to zero upon addition of the Fe catalyst, there was
no effect on the organic removal rate.
35
-------
Two samples of sewage (not shown) were run in the photo-
chemical apparatus at 60 g/1 zinc oxide. One sample consisted
of mixed industrial and domestic sewage and the other was domes-
tic sewage only. Due to heavy spring rains, we were unable to
obtain samples with appreciable soluble organic matter as a
result of rain-water dilution. The domestic sewage dropped from
an initial 20 mg/1 to 10 mg/1 organic carbon in 7 hours; the
mixed sewage dropped from an initial 33 mg/1 to 20 mg/1 in about
3 hours. Both samples leveled off, and further illumination for
3-4 hours had no effect. Both samples had been filtered through
fiber glass filter paper. The residual organic may have been
in colloidal form.
BACTERICIDAL PROPERTIES OP IRRADIATED ZINC OXIDE SLURRIES
A final experiment was conducted to determine the bacter-
icidal properties of illuminated zinc oxide. The test was con-
ducted in saline solution with E. coli under the Westinghouse
sunlamps in 2500 ml culture flasks in a low speed shaker.
Although the zinc oxide showed a higher kill rate at 30 minutes
and at 2 hours, the differences, compared with an illuminated
control, were not considered statistically significant. The
validity of the test was further diminished by the fact that
oxidation of phenol in the same apparatus was substantially less
than in crystallizing dishes with the same illumination and ZnO
concentration and time of exposure. The reason for this is not
clearly understood, sincePyrex transmits about 90% of the radia-
tion in this region (3000-4000A) . It may be that the angle of
incidence (variable from 0 to about 45° due to the contours of
the flask) resulted in a much higher reflectance than if it had
been normal to the light source, zinc oxide of both 1 g/1 and
10 g/1 ZnO had no inhibitory effect on the growth of E. coli in
the dark in nutrient.
While repeating the photo-oxidation tests with nutrient in
open containers was considered desirable, expiration of time and
funds did not permit this; thus, the question of the bactericidal
36
-------
and/or bacteriostatic properties of illuminated photocatalysts
is left in doubt.
37
-------
DISCUSSION
Although our original concept looked for the reaction of
oxidant with organic materials, it is probable that the mechanism
entails the direct reaction of an active oxygen species, atomic
oxygen or singlet excited state molecular oxygen. From Figures
3, 4, and 7, it is evident that beyond a certain level of organic
matter, the oxidant concentration has little relationship to
organic level. In the case of zinc oxide, Markham and Laidler
attributed this result to the fact that organic molecules are
adsorbed at the K2°2 decomP°sition sites, thus favoring the
formation reaction and inhibiting the decomposition reaction.
This apparently holds true for our oxidant as well (see Table 8) .
It must also be inferred that oxidant level is not a direct
measure of catalyst effectiveness. Although zinc oxide
apparently produces the highest oxidant level and apparently is
the most effective oxidizer (Table 7), the actual oxidation
data must be viewed with some reservation since they may repre-
sent merely differences in absorption of sunlamp irradiation of
TiO? vs. ZnO. The actual surface area of the pigment is also
an important factor.
We suggest that three conditions must be satisfied to
achieve oxidation of an organic molecule.
(a) the molecule must be adsorbed at, or be in
the vicinity of the active site on the catalyst.
(b) light energy of suitable wavelength (below about
4200A) must impinge on the active site.
(c) dissolved oxygen must be present to replace the
active oxygen species displaced by the radiation.
From these considerations, everything else being equal,
the higher the concentration of contaminant the faster the rate
of oxidation. Thus it would appear that photocatalytic oxidation
will find its greatest utility in problems of industrial waste
treatment where massive contamination is involved.
In our sunlamp exposures, we exceeded the manufacturer' s
recommended use period of 2000 hours by a factor of 2.5; however,
38
-------
not until the end of this period did a noticeable decrease in
oxidation rates appear. Extrapolation of the manufacturer's
life-intensity curves, shows that the tubes dropped to perhaps
20% of their original value before a decrease in efficiency
could be noted. From this, we infer that with this light source,
the limiting parameter is the number of active sites on the
catalyst surface and not the intensity of the source until it is
greatly attenuated.
Despite a very large number of tests, we find ourselves
without firm data on such a basic factor as the performance of
various zinc oxides. The selection of Photox 801 on the basis
of the H2O2 equilibrium was perhaps unfortunate; Photox 80 is a
much better electrophotographic material, but as Figure 11 indi-
cates, it is evidently not as good a photocatalyst.
Time did not permit a definitive investigation of tempera-
ture-dissolved oxygen relationships; however, we think the level
of dissolved oxygen is not as important as the rate at which dis-
solved oxygen is replaced. Thus, vigorous agitation at elevated
temperature would favor faster kinetics, despite lower dissolved
oxygen level.
The premise upon which the investigation was undertaken
has been amply confirmed. It is definitely possible to oxidize
a number of known organic materials as well as a substantial
portion of those in sewage by photocatalytic techniques. The
evidence that this process occurs in nature is convincing and
of great potential significance in pollution control.
39
-------
FUTURE WORK
Since it appears that with sunlamps, at least, the limiting
factor is the number of reactive sites on the catalyst particle,
examination of the wide variety of photocatalytic materials
available would be fruitful. A brief investigation under the
present program produced a substantial advance in TiO. perfor-
mance. The anatase used contained a rutile seed crystal, which
probably produces more crystallites in the anatase-rutile trans-
ition stage, with consequent greater photosensitivity. This
result might further be enhanced by water quenching from the
calcination temperature and/or by impurity doping.
An examination of the various grades of zinc oxide might
reveal much more active types than the ones we used. The exper-
iments with one type of beach sand suggest that a large number
of catalysts are available in natural form which might be iso-
lated and adapted to photocatalytic applications.
INTERACTION OF IONIZING RADIATION WITH METAL OXIDE CATALYSTS
This approach is based on work of Yamamoto and Oster who
studied the polymerization of calcium acrylate and acrylamide
in aqueous media under the influence of ionizing radiation.
These investigators reported an "enormous" increase in polymer-
ization effect when zinc oxide was added to the system. It
appears reasonable that photo-oxidation would show a similar
acceleration.
In such a system, a number of advantages over the photo-
chemical system would accrue. Most importantly, the radiation
would penetrate the slurry to a considerable depth so that much
less surface area would need to be exposed. Higher efficiency
should result because of the increased likelihood of the radia-
tion striking an active site. Finally, the energy cost, should
be much lower and the time of exposure much shorter.
40
-------
REFERENCES
1. Goodeve, C. F, and Kitchener, J. A., Trans Faraday Soc., 570
(1938) .
2. Jacobsen, A. E., I&E Chem., Vol. 41, No. 3, p. 523 (1940).
3. Weyl, W. A. and Forland, T., I&E Chem., 42, 257 (1950).
4. Bauer, E, and Neuweiler, C., Helv. Chim. Aota 10, 901 (1927).
\
5. Markham, M. C. and Laidler, K. J., J. Phys. Chem., 57, 362
(1963) .
6. Yamamoto, M. and Oster, G., J. Polymer Sci., Part A~l,
Vol. 4, 1683, 1688 (1966).
7. Pettijohn, F. J., Petrography of Beach Sands of southern
Lake Michigan, J Geology, 39, 432-455 (1931).
8. Oster G. and Yamamoto, M., J. Phys. Chem., 7C[ (10) 3033-6.
(1966) .
9. Bishop, D. F., Stern, G., Fleischman, M., and Marshall, L.S.,
I&E C Process Design and Development, Vol. 7, p.110 (1968).
10. Chambron, M. and Giraud, A., Academic Nationale De Medecine
Seance du 11, p. 631, Oct. 1960.
41
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APPENDIX
Analytical Procedures for the Determination of Low Concentrations
of Hydrogen Peroxide in Aqueous Solutions
.1. INTRODUCTION
The following outlines in detail the step-by-step proce-
dures to be followed in analyzing for hydrogen peroxide by the
potassium iodide method. None of the theoretical aspects will
be included in this description; however, the precautions to be
observed and special techniques to be followed will be described
in each step.
II. EQUIPMENT
1 - 5 ml micro burette 2 - glass stirring rods
1/100 ml graduations 1 - wash bottle distilled
2 - 1 ml pipette water
1 - 2 ml pipette 2 - 1 qt stoppered bottle
1 - 10 ml pipette 1 - 1 pt stoppered bottle
2 - 10 ml graduate Eye droppers
1 - 25 ml graduate 1 - acid dropping bottle
4 - 150 ml Erlenmeyer Flasks 1 - trip balance
1 - stop watch
Bausch and Lomb Spectrophotometer "Spectronic 20"
with constant voltage transformer.
III. CHEMICAL REAGENTS
1. Distilled water
2. 0.01 N standardized sodium thiosulfate (Na2S2O3«5H2O) .
Supplies of this solution may be obtained from chemical
supply houses. It should be kept in ordinary stoppered
bottles to avoid contamination or concentration changes.
(Dilute 0.1 N standardized solution to 0.01 N) .
3. Use H-SO, acid in KI to maintain acid pH.
4. 1 N potassium iodide solution: The potassium iodide (KI)
solution should be prepared daily by adding sufficient
distilled water to approximately 150 g of reagent grade
KI to make up 900 ml of solution. This solution must
Al
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be stored in an amber bottle and cannot be used if over
24 hours old.
5. Indicator: Starch indicator is prepared daily by adding
5 g of reagent grade soluble starch to 100 ml of boiling
distilled water and stirring vigorously. The solution
must be clear upon cooling. The indicator is useable
only on the day prepared.
Sampling Procedure
We assume a 10 ml sample is representative of the solution
we want to determine.
In a 20 ml beaker we place 10 ml of our unknown. Add 5
drops of 0.5 N H2SO4 - 20 drops of KI (in solution) - 5 drops of
starch solution.
The stop watch is started as the KI is added.
Pipette 1 ml of sample, add to 10 ml graduate, dilute to
10 ml label "A".
Pipette 1 ml of sample, add to 25 ml graduate, dilute to
20 ml ~ label "B".
Pipette 2 ml of sample, add to 10 ml graduate, dilute to
10 ml label "C".
After diluting, shake to evenly distribute solution.
Take B graduate, rinse the B & L test tube and fill to mark. Use
spectronic 20 - record transmission.
Take A graduate, rinse the B & L test tube back and forth, fill
to mark - record transmission.
Take C graduate, rinse the B & L test tube back and forth, fill
to mark - record transmission.
The dilutions and reading should be all done in a 5 minute
period. Total time on stopwatch should be under 10 minutes.
The transmission values are then used on Figure A-l to
determine the concentration of unknown solution.
A2
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35
30
25
20
a
a
0)
T3
H
X
o
Ul
c
(U
p
tc
15
10
Wavelength
488 490
00 90
1 1 1
80 70 60 50 40
1
30 20 1C
% Absorption
Figure A-l
ABSORPTION OF STARCH IODIDE
AT VARIOUS CONCENTRATIONS
A3
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Table A-l
ABSORPTION OF STARCH IODIDE AT VARIOUS CONCENTRATIONS
Sample
No.
A
B
C
D
E
P
G
B
I
J
Dilution
Value
3.29
ppm
6.58
9.87
13.16
16.45
19.74
23.03
26.32
29.61
32.90
Titration
Value
4.25
3.4*
7.47
6.62*
10.07
9.85*
13.2
12.4*
16.3
15.4*
19.9
19.0*
22.8
21.9*
25.8
25.0*
28.6
27.7*
32.4
31.6*
1/5 Dilution
Abs
92
71
54
43
34
29.5
19
18
14
11
1/10 Dilution
orption Values
97
93
86
81
75.5
70.5
61.0
59.0
54.0
47.0
1/20 Dilution
99-100
98
96
93
91
87
83.5
81
82
73
*Corrected for distilled water blank.
Dilution was made with distilled water.
10 ml tap water ~ 3 drop (0.14 ml Thiosulfate solution)
10 ml distilled water ~ I drop [o.05 ml Thiosulfate solution).
A4
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Titration Procedure can be used as a check.
Titrant 0.01 N Sodium Thiosulfate
20 ml of 1 N Potassium Iodide (KI)
20 drops of 0.5 N H-SCK
10 ml of sample solution Start stopwatch
5 ml of starch solution
Start titrating after 5 minutes of reactive digestion.
When solution starts going from purple > brown, add the
thiosulfate dropwise at a very slow rate until a sharp color
change from purple brown to water white is obtained. Wait 5
seconds between drops at this stage to avoid going past the end
point.
Amount of thiosulfate used is recorded.
Equations:
H2O2 + 2KI ) 12 + 2KOH
I2 + 2Na2S203 > 2NaI + Na2S4O6
Equivalent:
1 mole H2O2 1 mole I2 2 moles Na2S203
17 grams H~O2 1 mole Na2S2O3 (1 liter of 1 normal solution)
1 cc of 0.01 N Na2S203 = ^gog'01 = 0.00017 grains H,^
A5
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