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SILICA-TITANIA COMPOSITES FOR WATER TREATMENT
DANIELLE JULIA LONDEREE
A THESIS PRESENTED TO THE GRADUATE SCHOOL
OF THE UNIVERSITY OF FLORIDA IN PARTIAL FULFILLMENT
OF THE REQUIREMENTS FOR THE DEGREE OF
MASTER OF ENGINEERING
UNIVERSITY OF FLORIDA
Danielle Julia Londeree
First of all, there is no way I could have accomplished this research without the
help of my Lord and Savior. He has not only given me this opportunity, but also gave me
peace and patience during the times of frustration. Second, without the help of my
husband who kept me calm when my computer crashed several times, this work would
not be done. He also stayed up with me many nights while I was writing and constantly
gave me love and support.
I would like to thank Dr. David Mazyck for giving me the opportunity to work for
him and learn from him. He has challenged me in many ways and I feel that I have
learned more in the past two years of graduate school than the four years of undergrad. I
thank him for patience and for purchasing the spectrophotometer. I thank Dr. Powers for
allowing me to do an independent study with him. His excitement for research is
contagious and motivational. I would also like to thank Dr. Paul Chadik and Dr. Chang-
Yu Wu for giving advice during the research meetings and helping to provide answers for
my many questions.
I thank the students in my research group and for those who have helped me get
good data, specifically Matt Tennant (my mentor), Ameena Khan, Jennifer Hobbs,
Christina Termaath, Jon Powell, Noworat Coowanitwong, and Julee MacKenzie.
TABLE OF CONTENTS
A C K N O W L E D G M E N T S ......... .................................................................................... iii
LIST OF TABLES .............................. ......... .... ..... .. .... ....... ....... vi
L IST O F F IG U R E S .... ...... ................................................ .. .. ..... .............. vii
ABSTRACT ........ .............. ............. ...... .......... .......... ix
1 INTRODUCTION .................. .................................... ............ ..................
2 LITERA TURE REVIEW ............................................................... ............4
2.1 H eterogeneous Photocatalysis ........................... ....... .................................. 4
2.2 Titanium D dioxide .................. ................................ ....... .. .......... .. 8
2 .3 R actor D design ..................................................... 11
2.4 C analyst Supports ....................................................... .... ... ........ .. 13
2.5 Silica-Titania Composites ......... ................... ......... .. .... .............. 16
2.5.1 Sol-Gel Chem istry......................... ...... .. ....... ........ ....... 18
2.5.2 Surface Chemistry and Adsorption Characteristics .................................... 22
3 M ATERIALS AND M ETHODS ............................................................................25
3.1 Silica-Titania Com posites ......... .. ................. ................. ... ....................... 25
3.2 Batch Performance Studies .......... ........ ................. ... .............. 26
3.3 Colum n Perform ance Studies ........................................... .......................... 28
4 RESULTS AND DISCUSSION ............................................................................30
4.1 D ye Com prison ......... ..................... ... ........... ......... ........... 30
4.2 Optimization of Gels ........... ...... .. ..... .... .. .............. .......... 34
4.2.1 Titania Loading .................................... ........ ... .. .......... .. 34
4.2.2 Curing Temperature ........... .... ............... .................... 40
4.2.3 Pore Size ........ ....................... ...................... ........... 43
4.3 Column Studies ........................................................................ .... ......... .................. 44
5 SUM M ARY AN D CON CLU SION S ...........................................................................48
A COMMON PREPARATION METHODS OF MIXED AND SUPPORTED
OXIDES AS DISCUSSED BY GAO AND WACHS (1999). ....................................50
A 1 M ixed O xides ..................................................... 50
A 1.1 Sol-G el H ydrolysis.................................................. ......................... 50
A 1.2 C precipitation ............................................................ .. ........ ..... 50
A .2 Supported Oxides ............. .................. ......... .. ................ 51
A .2.1 Im pregnation .......... .. ........ ........................ ...... ............ ..... 51
A.2.2 Chemical Vapor Deposition ......................................................... 51
B D Y E ST R U C T U R E S .......................................................................... ....................52
B 1 M ethylene B lue ....................................... ............ ........... ....... ....... 52
B .2 M alachite G reen .................. ........................................ ............ .. 52
B .3 Crystal Violet ......................................... 53
B .4 R active R ed ........................................ 53
C T E M P IC T U R E S ..................................................................................................... 54
LIST OF REFERENCES ................................................................. ...........55
BIO GR A PH ICA L SK ETCH ........................................................................... 62
LIST OF TABLES
2-1 Oxidation power for various species relative to chlorine. ......................................6
2-2 Characteristics of Degussa P25 Titanium Dioxide............................. ...............
4-1 Peak absorbance w avelengths ................................................................... ....... 30
4-2 Exhaustion (C/Co = 1) times (minutes) for column studies. ................................46
LIST OF FIGURES
2-1 Schematic of oxidation and reduction occurring on the semiconductor surface. ....5
2-2 Example of hydroxyl radical attack a) on tetrachloroethylene by self-addition b)
on 1,1,1-trichloroethane by hydrogen abstraction (notice chlorine shift in the
ra d ic al) .............................................................................. 8
2-3 Diagram of anatase and rutile crystal structures..............................................10
2-4 Simplified diagram of a catalyst support. .................................... ...............13
2-5 Linkages of SiO 2 tetrahedras........................................................................ .... 17
2-6 Structural changes during the sol-gel process............................................ 20
2-7 Silanol groups on the silica surface ............................................ ............... 23
2-8 Representation of dehydroxylation............................................................23
3-1 C olum n setup. .......................................................................29
4-1 Photolysis of dyes as a function of UV flux for 2 hours of exposure .................31
4-2 Photolysis of MB and MG (2 hours UV exposure at 0.45 mW/cm2) as a function
of initial dye concentration. ....................................................................... ......... 32
4-3 Photolysis of RR (2 hours UV exposure) as a function of initial dye concentration
an d flu x ......................................................... ................ 3 2
4-4 Change in absorbance of RR at 538 nm as a function of pH .............................33
4-5 Spectrophotometer scan of 10 mg/L RR as a function of pH, before and after UV
4-6 Destruction ofRR (pH 7.6) after 2 hours UV exposure versus TiO2 loading as a
function of flux. ........... .. ............... .... .......... ................. ...... .....35
4-7 Destruction of RR (10 mg/L) after 2 hours UV exposure (0.45 mW/cm2) versus
TiO2 loading as a function of pH ................. ....... ....................................... 36
4-8 Destruction of RR (10 mg/L) after 1 hour UV with TiO2 slurry at various pHs ..37
4-9 Effects of bicarbonate at pH 7.6 versus TiO2 loading ............... .....................38
4-10 BET surface area and pore size versus titania loading. .......................................39
4-11 Adsorption of CV (10 mg/L) versus TiO2 loading after 24 hours of mixing........39
4-12 Destruction of RR (10 mg/L) after 2 hours UV exposure (0.45 mW/cm2) versus
curing temperature as a function of pH.......... ............. .......................40
4-13 Effect of curing temperature on surface area....................................... ........... 41
4-14 XRD analysis of various temperature-cured silica-titania composites.................42
4-15 Effect of curing temperature on adsorption of CV on gels with and without
T iO 2 ................................................................................ 4 3
4-16 Destruction of RR (10 mg/L) after 2 hours UV exposure (0.45 mW/cm2) versus
pore size. ...........................................................................44
4-17 Adsorption of CV on gels (6% TiO2) versus pore size. ......................................44
4-18 Column exhaustion curve for 140 A pellet (12% TiO2).....................................45
4-19 Column exhaustion curve for 30 A pellets (12% Ti2). .....................................46
4-20 Effect of regeneration time on 140 A column runs..............................................47
Abstract of Thesis Presented to the Graduate School
of the University of Florida in Partial Fulfillment of the
Requirements for the Degree of Master of Engineering
SILICA-TITANIA COMPOSITES FOR WATER TREATMENT
Danielle Julia Londeree
Chair: David Mazyck
Major Department: Environmental Engineering Sciences
Heterogeneous photocatalysis can be used for mineralizing organic dyes found in
the effluent of textile dyeing operations. Incorporating the catalyst TiO2 into an adsorbent
material, such as silica, has many advantages over using a TiO2 slurry for water
purification. The goal of this research was twofold: 1) to produce a silica-titania
composite using a sol-gel hydrolysis method that dopes the catalyst into the silica matrix
during gelation and 2) to optimize the titania loading, curing temperature, and pore size
of the material based upon maximizing its destruction and adsorption ability for textile
dyes. The optimal titania loading found for reactive red dye was 30 wt% TiO2. The
optimal pore size and curing temperature found were 140 A and 180C, respectively.
These composites were also made into small cylindrical pellets and tested in a flow-
through column to be used in a regenerative system. It was found that diffusivity was
very important to efficiently regenerate the column using photocatalysis.
The textile industry's dyeing operations are of primary environmental concern,
due to the variety of toxic chemicals (mostly solvents, surfactants, and acids) used in the
dyeing of garments. According to the Environmental Protection Agency [EPA] (1997),
chemicals such as methanol, methyl ethyl ketone, trichloroethylene, trimethylbenzene,
and dyes such as basic green 4 and dispersive yellow 3 are released into the environment,
either in the air phase or water phase. These contaminants can create large amounts of
BOD (biochemical oxygen demand) and COD (chemical oxygen demand), as well as
cause eutrophication and aquatic toxicity in effluent streams and natural water bodies.
Color itself is increasingly being regulated, as there have been links to aquatic toxicity
and lower DO (dissolved oxygen) values in receiving streams (EPA, 1996).
There are several chemical and physical methods of removal available for pre-
release treatment, such as adsorption by activated carbon, biodegradation, and ozonation.
Yet, dyes can be resistant to biodegradation and ozonation due to the large number of
aromatic rings found in modern dyes (Sheng and Chi, 1993). Therefore, advanced
oxidation processes are being considered as an emerging technology to handle large
volumes of textile wastewater (Sauer et al., 2002; Wu et al., 1999). Of these processes,
heterogeneous photocatalysis can be used for mineralization (complete conversion to CO2
and H20) of several classes of organic compounds (including dyes), instead of
transferring the pollutants from one phase to another (associated with carbon use).
Dyes can also be used as a surrogate for more toxic organic compounds in order
to measure the photocatalytic ability of a system. For instance, the structure of methylene
blue is similar to chloropromazine, an antipsychotic drug, and methylene blue destruction
could be used to estimate the destruction of this drug. In this research, dyes were used to
represent hard to degrade aromatic organic compounds in order to measure the
performance of the silica-titania composites for water treatment.
In addition, NASA, with its construction of the International Space Station and
plans for a Mars mission, requires a water recovery system that will minimize chemical
additions and energy and provide water at a potable level. The water source to be treated
will come mostly from grey water, including urine, humidity condensate, and wash water.
NASA Document CTSD-ADV-245 also lists numerous organic compounds (many
aromatic) that may be present in water in space flight (1998). The dyes used in this study
are surrogates for those compounds that may be present in space flight for NASA.
There were two main objectives of this research. First, to develop a silica-titania
composite by utilizing a sol-gel hydrolysis method and doping a highly efficient TiO2
catalyst into the silica matrix during gelation. These composites could also be created into
any shape and size, such as small cylindrical pellets to be used for a packed column.
Second, the titania loading, curing temperature, and pore size of the composites were to
be optimized based upon maximizing its destruction and adsorption ability for textile
dyes. The result would be a combined adsorbent/photocatalyst that efficiently adsorbs
contaminants into its pores for oxidation where the hydroxyl radicals would
predominantly exist. It was hypothesized that a packed column of these composites made
into small pellets could create a regenerable system that could be used over and over
2.1 Heterogeneous Photocatalysis
Advanced oxidation processes (AOPs) are considered to be an alternative water
treatment technology for removing harmful compounds and microbes from public water
supply, comparable with traditional practices such as activated carbon adsorption, air
stripping, and chlorine disinfection. AOPs can not only disinfect the water from virulent
microbes, but can also oxidize toxic heavy metals, organic pesticides and solvents,
chlorinated compounds, and inorganic chemicals from the water without producing a
waste stream unsuitable for disposal or consumption. For example, these processes can be
used for chemical spill cleanup, treatment of industrial effluents, and wastewater
treatment (Turchi and Ollis, 1990).
AOPs involve the generation of highly reactive oxidative species produced by
several pathways, including reactions with H202 and 03 with or without ultraviolet (UV)
irradiation and heterogeneous photocatalysis. Heterogeneous photocatalysis entails the
use of a solid photocatalyst (usually a semiconductor) in contact with either a liquid or
gas, while homogeneous photocatalysis uses a catalyst of the same phase as the
contaminated media. While UV use or an oxidant alone may yield partial degradation of
a compound, the combined use of UV with an oxidant (with or without a photocatalyst)
has been shown to yield complete mineralization of organic contaminants to carbon
dioxide (Ollis et al., 1991).
The reactants in heterogeneous photocatalysis in water are naturally occurring
species (such as oxygen, water molecules, and hydroxide ions (OH-)) as compared to
chemically supplied reactants (such as hydrogen peroxide and ozone) required for other
AOPs. When the surface of the photocatalyst absorbs a specific amount of energy, an
electron from the valence band jumps to the conduction band, thereby leaving a positive
"hole" in the valence band. These electrons (e-) and holes (h ) may either recombine,
releasing heat, or migrate to the surface of the catalyst. The reactive oxidative species are
then generated from reactions with the electrons or holes on the surface (see Figure 2-1).
hv Release of heat
Figure 2-1. Schematic of oxidation and reduction occurring on the semiconductor
The primary oxidizing species in water is the hydroxyl radical (*OH) (Turchi and
Ollis, 1990). It is the most reactive of the oxidizing agents (see Table 2-1) and can be
generated via several pathways:
H20 + h+ *OH + H+ (1)
OH- + h *OH (2)
02 + 2H+ + 2e- H202 (3a)
H202 + e *OH + OH- (3b)
H202 hL- 2*OH (3c)
02 + e 02- (4a)
02- + H202 *OH + OH- + 02 (4b)
As can be seen from equations 1 4b, molecular oxygen plays a dual role in
heterogeneous photocatalysis by 1) creating an oxidative species, the superoxide radical
(02-), and 2) acting as an electron acceptor to prevent the electrons and holes from
recombining. These functions are very important in establishing an efficient
photocatalytic system, for it has been shown that with an increase in oxygen
concentration, there is a subsequent increase in the reaction rate of contaminant
degradation (Gerischer and Heller, 1991). Similarly, hydrogen peroxide plays a parallel
Table 2-1. Oxidation power for various species relative to chlorine.
Species Relative Oxidation Power
Hydroxyl radical 2.06
Singlet oxygen radical 1.78
Hydrogen peroxide 1.31
Perhydroxyl radical 1.25
Chlorine dioxide 1.15
Source: Elizardo, 1991.
Adsorption of compounds to the surface of the catalyst is also important in an
efficient photocatalytic system. For instance, competition with other compounds can be
deleterious to the adsorption of specific contaminants on the catalyst surface. Chen et al.
(1997) showed that the effect of inhibition due to competitive adsorption will hinder the
degradation rate, especially in the presence of bicarbonate and phosphate. In addition,
carbonate ions will also scavenge the hydroxyl radicals, thereby decreasing the
destruction rate (Chen et al., 1997).
Another factor affecting adsorption is pH and the effect of pH is highly compound
specific. For example, in the photocatalysis of textile dyes, a lower pH enhanced
adsorption of Orange II (containing an anionic sulfonate group) to the catalyst surface,
thereby increasing the rate of mineralization (Wu et al., 1999), while an alkaline pH
increased adsorption and subsequent degradation of methylene blue (a cationic dye)
(Houas et al., 2001). Sauer et al. (2000) also found similar relationships between
adsorption and destruction of reactive dyes.
The photocatalytic mineralization of an organic compound begins with either
reactions with hydroxyl radicals (or other less significant oxidizing species) or by direct
oxidation with the holes on the surface of the catalyst. A hydroxyl radical "attacks" an
organic compound by either removing an available hydrogen atom to form water or
adding itself to any unsaturated carbon bonds (Grabner et al., 1991; Halmann, 1996;
Hoffman et al., 1995; Hoigne, 1990; Mao et al., 1991, 1992, 1993). Mao et al. (1991,
1992, 1993) found that the rate of oxidation of organic correlated with the C-H bond
strengths, proving that H atom abstraction by *OH may be a rate-limiting step.
Consequently, from these primary reactions, an organic radical is formed (see Figure 2-
2). The subsequent radical transformations and radical to radical interactions lead to an
array of intermediate products, eventually resulting in CO2 and H20. For chlorinated
compounds, HC1 is also produced.
Cl ci Ci CI
C Q= + *OH C C -OH
a) CI CI Cl
H Cl H Cl
I I I I
H-C-C-CI + *OH H-C-C. + H20
b) H CI ci Cl
Figure 2-2. Example of hydroxyl radical attack a) on tetrachloroethylene by self-addition
b) on 1,1,1-trichloroethane by hydrogen abstraction (notice chlorine shift in
the radical) (adapted from Mao et al., 1991 and 1992).
There are several literature reviews available on the mineralization of particular
classes of organic chemicals (Hoffman et al., 1995; Matthews, 1988; Ollis et al., 1991;
Serpone, 1995). Blake (1995) and Halmann (1996) have compiled extensive surveys of
this research and have also included sections on typical reactor designs and methods for
catalyst improvement. Research has also been focused on the determination of the
intermediate degradation pathways, for Mao et al. (1991 and 1992) has shown that the
conditions of a system (pH, oxygen concentration, competition with other compounds)
can affect the pathway a single compound can take towards complete mineralization.
2.2 Titanium Dioxide
Metal oxides are the most popular catalysts used for potable water and air
treatment; of these, TiO2 and ZnO have been the most researched and are considered the
most efficient. Of these two, TiO2 tends to be favored due to its stability in extreme
conditions (low or high pH), insolubility, heat resistance, non-toxicity, and low cost. TiO2
is also photostable as compared to ZnO which can undergo photocorrosion, decreasing
the effective lifetime of the catalyst (Okamoto et al., 1985).
Titanium is the ninth most abundant element in the earth's crust and occurs in
nature only in combination with other elements, such as with oxygen to form TiO2 (US
Geological Survey, 2002). Naturally occurring forms of titanium dioxide are usually
combined with iron (FeTiO3 or FeO-TiO2) and contain other impurities. Therefore, the
TiO2 (titania) used as a photocatalyst is manufactured from synthetic means.
Of the commercially available titanium dioxides, Degussa's P25 seems to be the
most researched for its photocatalytic ability in water (Halmann, 1996). Table 2-2 shows
a few specific characteristics of P25. The product is produced from high-temperature
flame hydrolysis of TiC14 in the presence of oxygen and hydrogen and is then steam
treated to remove HC1, resulting in a 70:30 anatase to rutile phase TiO2 (Degussa
Technical Bulletin, 1990). The phase of TiO2 is important since the anatase crystal form
is considered the more photoreactive of the two (Tanaka et al., 1991). Figure 2-3 shows
the crystal structures of the two phases. A difference in reactivity of the phases is
attributed to 1) the more positive conduction band of the rutile phase, hindering
molecular oxygen to act as an electron acceptor, and 2) the difference in surface
properties between the two phases (Tanaka et al., 1991).
Table 2-2. Characteristics of Degussa P25 Titanium Dioxide
BET Surface Area 50 m2/g
Average Primary Particle Size 21 nm
Band Gap: anatase, rutile 3.29 eV, 3.05 eV
Point of Zero Charge pH 6.0
Source: adapted from Degussa Technical Bulletin (1990)
A 11 skf ;^ !
Crystal morphology, as well as a variety of other factors, can influence the
k Y T4.5929A -9
I 4 -- C OXYGEN
Anatast Ru ile
Figure 2-3. Diagram of( anatase and rutle crystal structures (Millennium Chemical,
Crystal morphology, as well as a variety of other factors, can influence the
activity of a catalyst. Surface area is known to affect the adsorption ability of the catalyst,
as well as the available contact area for oxidation and reduction reactions. It would make
sense that a catalyst with a large surface area would have an increased rate of reaction,
but Tanaka et al. (1991) and Mills et al. (1994) showed that the destruction efficiencies of
several different titanium dioxide samples were independent of surface area (within the
range of 2.75 m2/g to 177 m2/g). Therefore, variables besides surface area may have a
more influential role in the activity of a catalyst. These variables are crystal morphology,
as well as crystallite size, prevention of electron-hole recombinations, zeta potential, and
band gap energy (Okamoto et al., 1985; Palmisano and Sclafani, 1997; Suri et al., 1993;
Tanaka et al, 1991; Zhang et al., 2001).
Tanaka et al. (1991) showed that the photocatalytic ability of a TiO2 catalyst is
foremost affected by the anatase content and secondarily affected by the size of the
crystal. For example, a photocatalyst containing larger crystals is more efficient than one
with smaller crystals, even when both contain equal percentages of anatase. This could be
due to the larger migration distance of the holes and electrons to the surface of the
catalyst, thereby decreasing the possibility of recombination (Tanaka et al., 1991).
Modifications in the catalyst surface have also been investigated for prevention of
electron-hole recombination. The addition of platinum and other transition metals have
been successfully arrayed on the titanium dioxide surface (Abrahams et al., 1985;
Okamoto et al., 1985; Martin et al., 1994; Suri et al., 1993). These metal additions have
an optimum at low weight percentages (less than 5%), above which the metal actually
hinders the photocatalytic ability.
The isoelectric point or point of zero charge (PZC) represents the pH at which an
immersed solid oxide would have zero net charge, resulting in electrically equivalent
concentrations of positive and negative complexes on the surface. At a pH above the
PZC, interactions with cationic electron donors and acceptors will be favored; while
anionic electron donors and acceptors will be favored at pH below the PZC.
The band gap is equal to the amount of energy required to activate the surface of
the catalyst. A band gap of 3.29eV corresponds to a wavelength of 378 nm. Therefore,
wavelengths below 378 nm (such as those emitted by UV light) would have the required
energy to activate the anatase phase of TiO2.
2.3 Reactor Design
A problem in the practical application of heterogeneous photocatalysis of
environmental pollutants is the design of a reactor that will maximize photocatalytic
efficiency while utilizing the least amount of energy. The simplest reactor configuration
uses an aqueous suspension or slurry of the photocatalyst. Most photocatalysis studies
use a slurry system in preliminary testing to establish the feasibility of pollutant
mineralization or microbe inactivation. For these systems, an optimum weight of TiO2
per volume of solution (wt/vol) was found that yielded the most degradation for specific
types of compounds at a specific light intensity. For instance, Goswami et al. (1993)
found 0.1% (wt/vol) TiO2 to be the optimum loading for most hydrocarbons and Block
and Goswami (1995) found 0.01% (wt/vol) as the best for microbial destruction.
Although slurry systems are considered to be the most photocatalytically efficient,
there is a dilemma to be addressed in the design of the reactor. Due to the extremely
small size of the particles (between 0.1 30 nm, depending on the manufacturing
source), there is a difficulty in recovering the catalyst from the purified water. This
problem has directed research to investigate the use of catalyst supports.
There are many research groups that have examined the immobilization of
photocatalysts on nonporous glass surfaces as a way of providing a backbone of support
for the particles, while also allowing the penetration of light to activate the catalyst
surface. Successful immobilization has been done on glass beads (Jackson et al., 1991;
Serpone et al., 1986; Zhang et al., 1994) and on the inside surface of glass tubes (Al-
Ekabi and Serpone, 1988; Matthews, 1988). The oxidation rates were usually lower with
the immobilized catalysts than with free suspensions. The determinate factor was
assumed to be the result of mass transfer limitations, for it was observed that as the flow
rate increased, resulting in more efficient mixing, the oxidation rate likewise increased.
While these designs may have certain advantages, it is hypothesized that an
efficient support for a catalyst would be a material with adsorption capability that would
bring the contaminants into close contact with the catalyst surface (see Figure 2-4) where
the hydroxyl radicals predominantly exist (Turchi and Ollis, 1990). Indeed, there has
been a reported increase in destruction efficiency when the reactants are brought into
close contact with the surface of activated TiO2 particles. This increase could be due to
the greater concentration of the substrate around the catalyst surface due to adsorption
interactions, for it has been proven that the highest efficiency of photocatalytic
degradation of organic pollutants is observed at relatively high concentrations of reagents
(Emeline et al., 2000).
Figure 2-4. Simplified diagram of a catalyst support.
2.4 Catalyst Supports
Adsorbents researched as possible catalyst supports include activated carbon,
silica gel, mordenite, alumina, zeolite, and aerogels (Takeda et al., 1995; Torimoto et al.,
1996; Yoneyama and Torimoto, 2000). The highest decomposition rate for propyzamide
(herbicide) was obtained with supports having medium adsorption strengths, such as
silica and mordenite (Torimoto et al., 1996). With supports of too low zeolitee) and too
high (activated carbon) adsorption strengths, no significant effect on the rate of
destruction was found.
Yet, research has focused on the benefits of using activated carbon (AC) as a
catalyst support versus other supports. Uchida et al. (1993) and Lu et al. (1999) found
that while the oxidation rates of a specific compound were lower with the TiO2 loaded
AC, the complete mineralization rates were greater than with a plain slurry of TiO2 or
any other catalyst loaded support. They concluded that the lower oxidation rate can be
attributed to the carbon blocking some of the photons of light from reaching the catalyst
surface and the higher mineralization rate can be attributed to the high adsorptive ability
of the carbon to adsorb the intermediate compounds that are being formed.
An additional benefit of TiO2 coated carbon is its potential for in situ AC
regeneration. Many water utilities use granular activated carbon (GAC) for the removal
of organic compounds. However, after the absorbent is spent, removal of the carbon from
the system is required for reactivation, disposal in a hazardous waste landfill, or
incineration. While on-site thermal reactivation is an option for some utilities, it is not the
most economical option for many. Therefore, research has been looking at optimizing
both TiO2-coated carbon's adsorption capability and its regeneration efficiency for the
removal and subsequent oxidation of organic pollutants adsorbed on the spent carbon
surface (Crittenden et al., 1993; Khan et al., 2002; Sheintuch and Matatov-Meytal, 1999).
While activated carbon has some benefits as a catalyst support, its use requires a
reactor that efficiently exposes the catalyst surface to the photons of light. Since this may
require mixing or fluidizing the particles, the inherent attrition can cause the catalyst to
detach itself from the carbon (Lu et al., 1999). An additional problem is that the surface
chemistry of a carbon may hinder effective coating (Khan et al, 2002). The various
methods of activating carbon can create different functional groups on the surface, thus
affecting the properties a specific carbon will display. This can greatly affect the
immobilization of a catalyst to the carbon surface.
Silica gels, on the other hand, have many advantages over activated carbon as a
catalyst support. The transparency of silica allows the penetration of photons to the
catalyst surface. This is extremely beneficial and allows for a fixed-bed reactor design
that can be highly efficient in relation to input energy. Silica also has high mechanical
strength, thermal stability, and can be synthetically formed into any shape, such as
cylindrical pellets (Yamazaki et al., 2001) or fibers (Brinker and Scherer, 1990). In
addition, silica is an adsorbent and is commonly used in chromatography columns for
adsorption and the resulting separation of compounds in an aqueous sample.
Furthermore, TiO2 and SiO2 can be chemically combined, enabling the formation of
highly efficient photocatalysts. These TiO2 SiO2 photocatalysts allow the placement of
the catalyst on both external surfaces and internal surfaces within the porous silica matrix
where pollutants are adsorbed.
Anderson and Bard (1995, 1997) found a strong synergy between silica and
titania as a combined oxide. For substances that are easily adsorbed to the silica surface, a
higher decomposition rate was found with TiO2 SiO2 composites versus a plain TiO2
slurry. The presence of an adsorbent was considered to promote efficiency by increasing
the concentration of the substrate near the TiO2 sites relative to the solution
concentration. But for substrates that are not readily adsorbed to the surface of silica, the
degradation rates were lower than compared to a slurry of pure TiO2. However, when
these rates were normalized to the TiO2 content of the material, the TiO2-SiO2 material
demonstrated a more efficient use of the TiO2 sites than for TiO2 alone. The highest
initial degradation rate was found with a 30/70 wt% TiO2-SiO2 composite.
Jung and Park (2000) also found 30 wt% titania (with gels prepared in a similar
process to Anderson and Bard) to be the optimum in their studies for the mineralization
of trichloroethylene. In addition, they concluded that the high porosity and large pore size
of the silica facilitated the mass transfer of reactants, resulting in a higher rate of
degradation than a plain slurry of Degussa P25.
Chun et al. (2001) used a different preparation method than Jung and Park (2000)
and Anderson and Bard (1995) and still found 30 wt% TiO2 to be an optimum. Chun et
al. additionally showed, using two organic compounds with differing characteristics, that
there is a strong correlation between adsorption of the compound on the mixed oxide
surface and the destruction rate of that compound. In conclusion, silica-titania composites
can be efficiently utilized in heterogeneous photocatalysis systems for the adsorption and
subsequent destruction of unwanted compounds in an aqueous solution. In addition, a
dosage of 30 wt% (equal to 12% on a weight per volume of silica precursor) was also
found to be an optimum loading in this research (see Chapter 4) for the silica-titanium
composites created using a sol-gel doping method (see Chapter 3).
2.5 Silica-Titania Composites
Silica is the most abundant oxide on the earth, yet despite this abundance, silica is
predominantly made by synthetic means for its use in technological applications. Since
synthetic silicas have a higher surface area than naturally occurring forms of silica, the
synthetic silicas provide the best adsorption and catalyst support structure for
heterogeneous photocatalysis. Although silica has a simple chemical formula (SiO2), it
can exist in a variety of forms, each with its own structural characteristics, as well as
chemical and physical properties. In general, SiO2 is a SiO4 tetrahedra, where each silica
atom is bonded to four oxygen atoms and each oxygen atom is bound to two silica atoms.
HO 0 OH
I I 1
Figure 2-5. Linkages of SiO2 tetrahedras (Hench and West, 1990).
There has been much research on the different preparation methods for creating a
solid SiO2-TiO2 material with photocatalytic ability. The characteristics (e.g., pore size,
surface charge, mechanical strength, and adsorption sites) of the final product are
dependent on the synthesis conditions and the type of interaction between TiO2 and SiO2.
There are two forms of interaction: physical forces of attraction (such as Van der Waals
forces) and chemical bonding (creation of a Ti O Si bond). The physically supported
TiO2 on SiO2 preparation methods have been the least researched, while the chemically
bonded TiO2 on SiO2 methods (also called mixed oxides) have been thoroughly
investigated (Gao and Wachs, 1999).
The most widely used methods of preparation (for details on these individual
methods, see Appendix A) for creating mixed oxides are sol-gel hydrolysis and
coprecipitation. Chemical vapor deposition, precipitation, and impregnation are the
methods commonly used for creating supported oxides, yet there is evidence of Ti-O-Si
bonds attaching the titania to the silica surface (Chun et al., 2001). A problem with
supported oxides is the inconsistency in the coating from one batch to the next and the
nonuniformity of the coating due to the sensitivity of the procedure to its conditions
(Hanprasopwattana et al., 1996; Lei et al., 1999). Another problem with any of these
methods is the assurance of the formation of an anatase phase of TiO2. Even though X-
ray diffraction (XRD) can be used to determine the phase of TiO2, the question remains
whether the preparation method has formed the most efficient semiconductor particle, as
compared to commercially available photocatalysts. For example, Liu and Cheng (1995)
produced a TiO2-SiO2 mixed oxide using coprecipitation and found that the SiO2 inhibits
the growth of TiO2 crystals, resulting in amorphous TiO2 within the silica matrix.
A method of titania incorporation in silica has been devised to overcome both of
these potential problems. This method uses sol-gel hydrolysis to create the silica matrix,
then during gelation the solution is doped with the commercially available highly
efficient Degussa P25 TiO2. The silica network may either form around the titania
particle and/or form a bond with the titania to secure its position in the matrix.
2.5.1 Sol-Gel Chemistry
A sol is a colloidal suspension of particles, while the term gel refers to the semi-
rigid material formed when the colloidal particles link together in a liquid to form a
network. Sol-gel precursors consist of metals or metalloids surrounded by various
ligands, of which the most widely used are metal alkoxides (Mauritz, 2002). Aluminates,
titanates, borates, and silicates are just some of the alkoxides that can be used to create a
sol-gel. A gel that contains SiO2 networks is given the term silica gel.
Hydrolysis and condensation. Gels derived from alkoxides form through
hydrolysis reactions followed by condensation reactions (Equations 5-7). Generally
speaking, hydrolysis replaces the alkoxide group with a hydrogen ion while condensation
produces siloxane bonding (Si-O-Si) with the products of the reaction being water or
alcohol. As the number of siloxane bonds increases, the individual molecules join
together to create a silica network. See Equations 5-7 (Mauritz, 2000).
-Si-OR + HOH -Si-OH + ROH (5)
-Si-OH + -Si-OH i -Si- + HOH
-Si-OH + -Si-OR --Si-O-Si- + ROH (7)
Gelation. After condensation reactions begin, gelation can occur (see Figure 2-6
for structure changes). Gelation is considered the growth period for the colloidal
particles, which grow by polymerization or aggregation, depending on the conditions.
Eventually, as the last links are made, the sol's viscosity increases to the point where the
sol becomes a semi-rigid gel that consists of a network of pores filled with liquid.
Aging. Aging occurs after gelation. Here, terminal silanols re-orient themselves
and react with each other to form additional network linkages. This results in shrinkage
of the gel and forces the removal of some liquid from the pores. As the gel stiffens,
"Ostwald ripening" occurs. This is the phenomenon whereby necks begin to form
between primary particles causing the filling of small pores (see Figure 2-6 part c).
A UQUID B LIOUDQUD NECKING
SQL GEL AGED GEL
NO 0 PORE
DRIED GEL PARTIALLY
Figure 2-6. Structural changes during the sol-gel process (adapted from Hench and West,
Drying. The drying step involves the removal of the solvent from within the pores
of the gel. This is very important in the determination of pore size. There are three types
of gels characterized by the solvent removal process: xerogels, aerogels, and cryogels. As
defined by Legrand (1998), a xerogel is obtained by evaporation of the liquid component
at ambient pressure and temperatures below the critical temperature of the liquid; an
aerogel is obtained from the evaporation of the liquid component above its critical
temperature; and a cryogel is obtained from the freezing and subsequent sublimation of
the liquid component. Aerogels have higher surface areas than xerogels and do not have a
material volume loss due to shrinkage. Yet, because supercritically drying gels can be
cumbrous and expensive, research has focused on making xerogels with similar
characteristics to aerogels. As far as water treatment is concerned, xerogels would be
preferred over aerogels due to the higher mechanical strength of the xerogel.
Curing. Calcination or curing is the final heat treatment step in producing a
mechanically strong gel. This thermal treatment also affects several characteristics of the
silica gel. With increased heat treatment, the BET (Brunauer, Emmett, and Teller
equation) surface area will decrease due to sintering (Papirer, 2000) and the dehydration
(around 180-200 C) and subsequent dehydroxylation (above 2000 C) of the surface of
the silica, resulting in the formation of siloxane bridges (Holysz, 1998). The final product
resulting from all these steps hydrolysiss, condensation, gelation, aging, drying, curing) is
a gel containing a relatively monodisperse pore size and displaying specific
characteristics associated with the conditions it experienced during processing.
Effect of catalyst. The two alkoxides most often used are tetramethoxysilane
(TMOS) and tetraethoxysilane (TEOS). TMOS has the advantage of rapid hydrolysis
under a variety of conditions, but the toxic methanol produced can be hazardous to the
eyes and lungs. Conversely, TEOS is hydrolysis rate limited (Brinker and Scherer, 1990),
but produces less toxic ethanol during the reaction. To increase the hydrolysis reaction
rate of TEOS, an acid or base can be added as a catalyst. However, the nature and
concentration of the catalyst will affect the characteristics of the gel. For example, the
addition of hydrofluoric acid will rapidly increase the rate of gelation as well as increase
the pore size (Powers, 1998).
Other important parameters to consider in the production of a gel are H20 to Si
molar ratio (called the "R" factor) and the aging and drying schedules. A high R ratio will
promote more rapid hydrolysis, for the concentration of water affects the rate of
hydrolysis. The drying time and temperature also affect the pore size of the gel as well as
the shape of the gel. For example, the longer the time to dry, the more Ostwald ripening
occurs and the less likely the gel will crack, thereby producing a larger pore size than a
gel of the same composition dried in short period of time (less than 48 hours). In
summary, there are many parameters in the production of a gel that will greatly affect the
final product's characteristics. It is these differences in silicas produced by various
preparation methods that make the comparison of data found in the literature a very
2.5.2 Surface Chemistry and Adsorption Characteristics
The surface of silica consists of two types of functional groups: silanol groups (Si
- O H) and siloxane groups (Si O Si). The silanol groups are the locale of activity
for any process taking place on the surface, while siloxane sites are considered
nonreactive (Unger, 1979). Porous amorphous silica contains three types of silanols on its
surface: isolated, geminal, and vicinal (or associated). Figure 2-7 shows the distinction
between these groups. The unequal distribution of the silanols in the sililca matrix,
resulting from irregular packing of macromolecules as well as incomplete condensation,
results in a heterogeneous surface (non-uniformity in the dispersion of silanol groups) for
synthesized silica (El Shafei, 2000).
The various silanols can have different adsorption activity and current knowledge
indicates that the isolated silanols are the more reactive species (Nawrocki, 1997; El
Shafei, 2000). With increasing temperature of heat treatment, the silica surface becomes
hydrophobic due to the condensation of surface hydroxyls (dehydroxylation) and the
formation of siloxane bridges (see Figure 2-8). This increases the ratio of isolated sites
versus other silanol groups on the surface and therefore increases adsorption.
geminal isolated single associated silanols
H H H
hydrogen bounded water molecules
Figure 2-7. Silanol groups on the silica surface (Legrand, 1998).
At temperatures up to 450C, the removal of the hydroxyl groups does not result
in visible changes in structure, density, or specific surface area of the silica because the
rehydroxylation makes it possible to return to the initial state (Davydov, 2000).
Rehydroxylation of the silica surface by exposure to water vapor is reversible up to
600C. At temperatures above this, the rehydroxylation becomes a slow process or
remains incomplete (Unger et al., 2000). Also, there is a decrease in surface area as a
result of sintering at temperatures above 500C (Persello, 2000).
H- Oi H- -...-H
Si Si + HO
Figure 2-8. Representation of dehydroxylation (Burneau and Gallas, 1998).
Pore size can be affected by the heat treatment of the silica, but is also a function
of the manufacturing process. The effect of pore size on adsorption is dependent on the
compound's characteristics, such as molecule size, shape, and functional groups
(Dabrowski, 2001; Goworek et al., 1997; Goworek et al., 1999). In general, the smaller
the pore, the larger the surface area and the larger the capacity for adsorption of
contaminants. Yet, Goworek et al. (1997) found that some compounds may be hindered
by a narrow pore size due to steric forces. Therefore, the adsorption ability for that
particular compound would be higher in a silica with larger pores.
Adsorption on silica surfaces is primarily due to ion interactions and hydrogen
bonding with the silanol groups. Therefore, the adsorption kinetics in relation to these
silanols is greatly dependent on the pH. Silica's point of zero charge ranges between 2
and 3 (Papirer, 2000; Persello, 2000). At a pH less than the PZC, the surface has a
positive charge due to the formation of Si-OH2 At a pH above the PZC, the surface of
the silica has a negative charge due to the deprotonation of the silanol group resulting in
Si-O-. As the pH approaches 7, the Si-O- sites become significant (the pKa for SiOH =
SiO- + H is 7) (Cox, 1993). These Si-O- sites readily react with cations in solution and
an increase in pH would greatly increase the adsorption of these compounds (Kinniburgh
and Jackson, 1981). For nonionic compounds in solution, adsorption can occur at the
silanol groups by hydrogen bonding or by Van der Waals forces. For anionic compounds,
there may be some adsorption at a pH above the PZC of silica, but it would be occurring
at the TiO2 surface because it is positively charged at pH < 6.0. One factor to consider
when adjusting pH in order to achieve the highest adsorption onto silica-titania
composites is the dissolution of silica at a pH greater than 9 (Iler, 1979). In order to
create a long-lasting photocatalytic composite, the dissolution of silica would need to be
MATERIALS AND METHODS
3.1 Silica-Titania Composites
The silica-titania composites were made by a sol-gel method (Powers, 1998)
using nitric acid and hydrofluoric acid as catalysts to increase the hydrolysis and
condensation rates, thereby decreasing the gelation time. The basic formula used to create
gels with a pore size of roughly 140 D is as follows: 25 mL water, 50 mL ethanol, 35 mL
TEOS (tetraethylorthsilicate), 4 mL nitric acid (IN), and 4 mL HF (3%). TEOS was used
in place of TMOS (tetramethylorthosilicate) because it produces less toxic ethanol
instead of highly toxic methanol that would be created by reactions with TMOS. Varying
the volume of hydrofluoric acid allowed the manipulation of the pore size. Decreasing the
volume to 1 mL HF resulted in 30 A pores, while increasing it to 8 mL resulted in 320 A
The chemicals (reagent grade from Fisher) were added individually, in no
particular order, to a polymethylpentene container. A magnetic stir plate provided
sufficient mixing. The solution was allowed to mix until gelation occurred (2 hours for
140 A gel). During this time, a known mass of Degussa P25 was added to the batch and
the percentage of titania recorded is given as a percent by volume of silica precursor.
After gelation, the lid was put on, preventing any air exposure in order to protect against
premature evaporation. Next, it was aged at room temperature for two days, then at 650 C
for two days. After aging, the gel was removed from the container, rinsed with deionized
water to remove any residual acid or ethanol, and placed in a Teflon container for the
next series of heat treatments. A small hole in the lid of the container allowed slow and
uniform drying of the gel. It was then placed in an oven and the temperature was ramped
from room temperature to 1030 C (2/min) and kept constant for 18 hours, resulting in the
vaporization of liquid solution within the silica network. Next, the temperature was
ramped to 1800 C (2/min) to remove any physically adsorbed water. It was kept constant
for 6 hours and then was slowly decreased back to room temperature over a 90 minute
period. These gels were powdered and sieved through 325 mesh to be used for batch
The titania-doped pellets were made in a similar fashion, except the solution
(including the P25) was pipeted into polystyrene 96-well assay plates before gelation.
The volume added to each well was approximately 0.3 mL. The plates were then covered
with lids and wrapped in foil to prevent premature evaporation. The same aging and
drying schedule as previously discussed was used, but before the drying step, the pellets
were removed from their plates and placed in a Teflon container. The resultant size of an
individual cylindrical pellet after drying was approximately 5 mm in length with a
diameter of 3 mm. These pellets were used in column performance studies.
3.2 Batch Performance Studies
The batch performance studies were done to determine an optimal titania loading,
curing temperature, and pore size that is efficient for both adsorption and destruction of
contaminants. The powdered gels were compared based upon destruction (removal of
color) of reactive red dye (RR) from Sigma-Aldrich and adsorption of crystal violet dye
(CV) from Fisher. All dye solutions were made with nanopure water. The destruction
studies were performed as follows: 30 mg of gel was added to 100 mL of a 10 mg/L RR
solution buffered with 200 mg/L of sodium bicarbonate in a 125 mL Erlenmeyer flask.
Then pH adjustments were made with sodium hydroxide or nitric acid. The flask opening
was then covered with parafilm to maintain a closed system with a limited amount of
dissolved oxygen (DO). In a system designed for NASA, oxygen would be a limiting
factor due to the problem of continually supplying oxygen for water treatment for long
duration space flight. The flasks were placed within a box containing four magnetic
mixers relatively equidistant from a single 4-Watt UV lamp. The UV exposure time was
two hours and the percent destruction was calculated by comparing the initial and final
absorbance readings using a Hach DR/4000U (Loveland, Colorado) Spectrophotometer at
RR's peak absorption wavelength of 538 nm. The spectrophotometer gave a linear
response across the concentration range for each dye. The UV intensity (mW/ cm2) was
varied by changing the position of the mixers in regard to the lamp and by the use of
additional lamps. UV intensity (wavelengths between 320 to 380 nm) was measured with
a VWR Ultraviolet Light Meter (Atlanta, GA). Each mixer was placed where the desired
UV intensity was achieved at the mixer's center.
The adsorption studies were performed as follows: 30 mg of gel was added to 100
mL of 10 mg/L CV solution with no bicarbonate added. The studies were conducted at a
pH of 4, using nitric acid and sodium hydroxide for the pH adjustments. The flasks were
covered with parafilm and placed in a box with no penetration of light and allowed to mix
for 24 hours. The percent adsorbed was calculated from the absorbance reading at a
wavelength of 590 nm.
The BET (Brunauer, Emmett, and Teller equation) surface area and pore volume
analysis was performed on a Quantachrome NOVA 1200 Gas Sorption Analyzer
(Boynton Beach, FL). The powdered samples were outgassed at 110 C for
approximately 24 hours and analyzed using nitrogen adsorption isotherms. The average
pore size was calculated by dividing the pore volume by the surface area. Pore size
distribution curves were also attained and coincided with the calculated pore size.
XRD (X-ray diffraction) analysis was performed on cured gels to determine the
phase of titania, anatase or rutile. These gels were cured at temperatures up to 8000 C as
an additional curing step after the regular drying schedule. The analysis was performed
on a APD 3720 from Philips Analytical (Almelo, Netherlands). The reflections of the two
crystalline TiO2 phases, anatase (20 = 25.3) and rutile (20 = 27.4), were taken from
Grieken et al. (2002).
3.3 Column Performance Studies
The purpose of these studies was to show the performance of the pellets in a
regenerative system, which places equal significance on the adsorption of a contaminant
as well as the feasibility of the contaminant to be mineralized. A cylindrical column (6
inches in length, 0.5 inch diameter) with a porous frit was filled with 10 mL of pellets.
The column setup is shown in Figure 3-1. A CV dye solution of 2 mg/L was pumped
through the column until the effluent concentration equaled the influent concentration.
Then the CV pump was turned off and the deionized water and UV light was turned on
(flux = 1.1 mW/cm2) for a set time (e.g., one hour). During this regeneration time, the
water was pumped through the column to provide the reactants (02, H20) required for
photocatalysis. After the regeneration time, the CV pump was turned back on. This cycle
was continued several times for separate packed columns of 140 A gels and 30 A pellets,
both loaded with 12% TiO2. The flowrate for both the deionized water and CV dye was 8
Figure 3-1. Column setup.
RESULTS AND DISCUSSION
4.1 Dye Comparison
In order to accurately measure the photocatalytic ability of the silica-titania (Si02
- TiO2) composites, it was important to choose a dye that displayed little or no photolytic
behavior, i.e. change in absorbance as a response to UV light exposure alone. The
following four dyes were examined in order to determine the most suitable dye for the
studies herein: malachite green (MG), methylene blue (MB), crystal violet (CV), and
reactive red (RR). See Appendix B for structures of these dyes.
Using a spectrophotometer, each dye was scanned at a range of 290 nm to 800 nm
in order to find each dye's peak absorbance wavelength. A summary of the peak
wavelengths is shown in Table 4-1. The absorbance measured at this peak wavelength
was used for comparing initial and final readings.
Table 4-1. Peak absorbance wavelengths.
Dye Peak Wavelength
MG 619 nm
MB 656 nm
CV 590 nm
RR 538 nm
A 100 mL solution of each dye (10 mg/L) was prepared in a 125 mL Erlenmeyer
flask and exposed to three different intensities of UV light for two hours. The findings are
presented in Figure 4-1. Each dye experienced a change in absorbance of less than 5% at
each flux, except for crystal violet, which experienced a 19% change at a flux of 1.1
mW/cm2. Based upon this undesirable change, crystal violet was excluded for future
Subsequently, the remaining dyes were evaluated for photolysis at various
concentrations (0.5 mg/L to 10 mg/L). At 0.5 mg/L, there was a 98% and 20% change
between the initial absorbance readings for MB and MG respectively, compared to their
final readings after two hours of UV exposure (0.45 mW/cm2) (Figure 4-2). Large
discrepancies were also observed at 1 mg/L and 2 mg/L. Therefore, MB and MG were
eliminated as viable candidates to investigate the photocatalytic ability of the composites.
For RR (see Figure 4-3), the largest change in absorbance was only 13% at a flux of 1.1
mW/cm2, the highest flux achievable. Thus, RR was chosen for the photocatalysis
studies. The amount of photolysis was not subtracted from the results of the destruction
c 0.45 mW/cm^2
S20 -- 0.75 mW/cm^2
] 1.1 mW/cm^2
RR CV MB MG
Figure 4-1. Photolysis of dyes as a function of UV flux for 2 hours of exposure.
- MB Initial Absorbance
* MG Initial Absorbance
---.--- MB Final Absorbance
MG Final Absorbance
'S 1.4 ---
ns = 1
S 0.8 98% for MB
S0. and 20% for
$ E 0.6 -----
0 0.4 M,
45% for MG
< ^B Y and 20% for
0 2 4 6 8 10
Figure 4-2. Photolysis of MB and MG (2 hours UV exposure at 0.45 mW/cm2) as a
function of initial dye concentration.
-o- Initial Absorbance
- Final Absorbance (1.1 flux)
---- Final Absorbance (0.45 flux)
0 1 2 3 4 5 6 7 8
9 10 11 12
Figure 4-3. Photolysis of RR (2 hours UV exposure) as a function of initial dye
concentration and flux.
Besides photolysis, pH can also affect a dye's peak absorption wavelength.
Therefore, RR solutions of 10 mg/L were prepared at different pHs and compared before
and after UV exposure. The change in absorbance (at 538 nm) as a function of pH is
shown in Figure 4-4. At a pH above 8, there was a decrease in absorbance along with a
slight visible change of color from a dark pink to a red. Figure 4-5 displays a section of
the spectrophotometer scan and shows the changes that occur at an alkaline pH. At a pH
of 7, the maximum peak wavelength is 538 nm. At a pH of 9.5, the highest peak occurs at
both 514 nm, as well as 538 nm. After two hours of UV exposure, the max peak returns
back to 538 nm. Also noticeable at pH 9.5 is a slight downshift in the curve after UV
exposure. Therefore, the pH was kept below 8 for comparison studies of gels to ensure
accurate measurements of destruction ability.
S0.1 -- Abs. (initial)
0.05 -- Abs. (final)
0 2 4 pH 6 8 10 12
Figure 4-4. Change in absorbance of RR at 538 nm as a function of pH.
Figure 4-4. Change in absorbance of RR at 538 nm as a function ofpH.
Absorbance pH 7 ------.Absorbance pH 9.5
-- Absorbance ph 9.5 after UV
400 450 500 550 600
Figure 4-5. Spectrophotometer scan of 10 mg/L RR as a function of pH, before and after
4.2 Optimization of Gels
4.2.1 Titania Loading
In order to determine the optimal titania loading, 14 gels (140 A) of varied TiO2
loadings (from 0.77% to 40%) were created. The desired optimal loading would be found
as the composite having the highest rate of destruction for RR. These gels were
performance tested at three UV intensities for the destruction of RR (10 mg/L) and the
results are displayed in Figure 4-6. There is a definite plateau where the percent
destruction remains constant, even with the addition of more TiO2. It appears that this
plateau occurs at lower TiO2 percentages with higher UV intensities. For 1.1 mW/cm2,
the plateau occurs at 5% TiO2 and it occurs at 10% for 0.75 mW/cm2. At the lowest UV
intensity, the destruction levels off at 12%. These results were triplicated and the error
bars are included in the figure.
2 A 0.75
0 5 10 15 20 25 30 35 40
Figure 4-6. Destruction ofRR (pH 7.6) after 2 hours UV exposure versus TiO2 loading as
a function of flux.
The plateau can be explained by the TiO2 is agglomerating in solution during gel
preparation, which was visually observed with TiO2 loadings greater than 8%, and
therefore is decreasing the effective surface area of titania available for excitation and/or
oxidation and reduction reactions on its surface. The 12% optimal loading (wt/vol
precursor) at the lowest energy is approximately equal to 30 wt%, which was found to be
an optimum for destruction of organic compounds with other methods of producing
silica-titania composites (Jung and Park, 2000; Chun et al., 2001; Anderson and Bard,
The structure of RR includes sodium sulfonate groups that would make the
compound anionic when the sodium disassociates as the compound dissolves. Therefore,
the effect of pH on adsorption and destruction of RR was also investigated. It was
observed that at a pH range of 3 to 10, there was no adsorption of RR dye to the silica-
titania gels. Thus, the color disappearance of the dye can be directly attributed to
photocatalysis with a minimal amount due to photolysis.
The effect of pH on the destruction of RR is shown in Figure 4-7. It was observed
that acidic conditions enhanced the photocatalytic destruction of RR and it was initially
hypothesized that this was due to the TiO2 having a positive charge at pH 3, therefore, the
amount of RR adsorption that would occur on pure TiO2 at low pH would increase the
pH 3, bicarb m pH 4, no bicarb
A pH 7.6, bicarb pH 7.6, no bicarb
0 1 2 3 4 5 6 7 8 9 10 11 12 13
Figure 4-7. Destruction of RR (10 mg/L) after 2 hours UV exposure (0.45 mW/cm2)
versus TiO2 loading as a function of pH.
Yet, when a slurry of titania was tested for destruction and adsorption at a range
of pHs, the results were contrary to what was hypothesized (Figure 4-8). While the
adsorption onto the titania surface may have slightly increased at a lower pH, the level of
total removal (destruction and adsorption) did not change. It is then presumable that at a
pH greater than 6, the silica surface becomes more concentrated with SiO- groups, thus
repelling RR. As the pH approaches 3, the silica has less of a negative charge and would
allow the penetration of RR into the pores. Hence, acidic conditions would increase the
ability of RR to get within the vicinity of the titania.
Destruction U Adsorption
70 70 c
0 60 60
0 50 50 C.
S 40 40 0
30 30 "
0 1 2 3 H 4 5 6 7 8
Figure 4-8. Destruction of RR (10 mg/L) after 1 hour UV with TiO2 slurry at various
Grey water contains many sulfates, phosphates, and salts that may inhibit the
destruction of unwanted compounds (Chen et al., 1997). The addition of sodium
bicarbonate allows one to see the performance of the silica-titania composites in a
realistic feed stream. In Figure 4-7, it is also interesting to notice the affect of bicarbonate
on the destruction efficiency. There is little increase in destruction (with no bicarbonate)
at a pH of 7.6 within a range of 0.77% to 6% TiO2. At a larger TiO2 percentage, the
difference becomes very noticeable. However, Figure 4-9 shows that the increase in
destruction without bicarbonate present is only seen within a certain range of TiO2
loading (8% to 12%).
no bicarbonate Awith bicarbonate
S 0 -
5 30 -
0 5 10 15 20 25 30 35 40
Figure 4-9. Effects of bicarbonate at pH 7.6 versus TiO2 loading.
It was also thought that surface area may play a role in the leveling-off seen in
destruction of RR. As can be seen in Figure 4-10, there is no real trend between BET
surface area and titania loading, as well as between pore volume and titania loading.
Notice that this surface area is normalized on a volume basis. This trend is different for
silica-titania composites prepared by other methods, for which there is a drastic change
(up to a loss of 200 m2/g) in surface area with an increase in percentage of titania
(Greiken et al., 2002). In conclusion, there is no correlation between surface area or pore
volume and the destruction of RR.
For adsorption studies, CV dye was used because of its cationic nature resulting in
a high affinity for the silica-titania composites. These studies were conducted at a low pH
(worst case scenario for adsorption of a cation contaminant onto a silica surface). Figure
4-11 compares the amount of adsorption versus titania loading. There is an increase in
adsorption with additional titania in the gel. It is interesting to note, however, that
adsorption, similar to the destruction studies, reached a plateau at 12% loading.
Therefore, based upon the destruction studies and adsorption studies, an optimal titania
loading of 12% TiO2 was chosen.
BET SA/vol A Pore Volume
E 40 *
30 AAAA A 0.2a
AA A A A A 0
j 20 A
0 5 10 15 20 25 30 35 40 45
Figure 4-10. BET surface area and pore size versus titania loading.
0 5 10 15 20 25 30 35 40 45
Figure 4-11. Adsorption of CV (10 mg/L) versus TiO2 loading after 24 hours of mixing.
4.2.2 Curing Temperature
The effect of curing temperature was investigated on gels doped with 6% TiO2
and those with no TiO2 (all made with the 140 A initial formula). The literature showed
that there can be an increase in adsorption with gels produced at higher curing
temperatures (up to 4000C) due to increasing the ratio of isolated silanols on the surface
versus other types of silanols. Yet, the curing temperature may have detrimental effects
on the photocatalytic ability of the composites. In looking at the destruction of RR (10
mg/L), Figure 4-12 shows that there is a 30% difference between the 180C gel versus the
500C gel at a pH of 7.8. This difference is less pronounced at pH 4. At 800C, there is a
marked decrease in destruction, regardless of pH. This can be explained by both a
decrease in surface area of the gels at 800C (see Figure 4-13) from an average of 275
m2/g to 150 m2/g for the titania loaded gel and a phase change of titania.
*pH 4 m pH 7.8
0 100 200 300 400 500 600 700 800
Curing Temp (oC)
Figure 4-12. Destruction of RR (10 mg/L) after 2 hours UV exposure (0.45 mW/cm2)
versus curing temperature as a function of pH.
*6% 102 0% 102
0 200 400 600 800
Curing Temp (OC)
Figure 4-13. Effect of curing temperature on surface area.
The anatase phase of TiO2 begins transforming into rutile phase at a temperature
of 600C and becomes complete at 1000C (Zhang et al., 2001). It was thought that the
shift in the phase of titania at 800C may have also affected its photocatalytic ability, but
XRD analysis (Figure 4-14) revealed only a slight shift from the anatase phase (29 =
25.3) to the rutile phase (29 = 27.4) at 800C which may or may not be enough to have
caused a drastic decrease in destruction. It may also be that the phase change is mostly
occurring on the surface of the titania where the reactions take place. In other words, the
large anatase peak is due to internal anatase crystal formation and the rutile formation is
occurring foremost on the surface where it is greatly affecting photocatalysis.
ze ......... *2" ......... 4* ..*.. *.... ** *'* ** ... ... ** ^.[' 3 e
20 22 24 26 28 [201 38
Figure 4-14. XRD analysis of various temperature-cured silica-titania composites.
In looking at adsorption of CV at various curing temperatures, Figure 4-15 shows
800C with the highest adsorption capacity even though it has the lowest surface area.
This may be explained by the cleaner surface (decrease of silanol sights) at this
temperature, resulting in the gel displaying a more hydrophobic behavior (Holysz, 1998).
This would decrease the competition between water and CV, allowing CV to efficiently
interact with the isolated silanols. However, the clean surface can slowly become
rehydroxylated when it is in contact with water (Davydov, 2000; Unger et al., 2000). The
adsorption ability would then most likely follow the same trend as gels that were treated
at lower curing temperatures. Therefore, based upon this and the destruction studies, the
optimal curing temperature chosen was 180C.
6% Ti02 0% Ti02
0 100 200 300 400 500 600 700 800 900
Curing Temp (oC)
Figure 4-15. Effect of curing temperature on adsorption of CV on gels with and without
4.2.3 Pore Size
Pore size was also investigated for its effects on destruction ability, as well as
adsorption ability. Figure 4-16 shows that a large pore size (>140 A) increases the
destruction rates for RR. This is consistent with what is found in the literature. Jung and
Park (2000) concluded that high porosity and large pore size facilitate the mass transfer
of reactants, such as oxygen and reaction intermediates. Thus, the increase in surface area
by small pores is not always effective for high photoactivity. The opposite is true for
adsorption, because a smaller pore size results in more surface area available for
adsorption. Figure 4-17 shows the largest adsorption with a pore size of 30 A. Adsorption
decreases 73% with larger pores (> 30 A), and in this range there is no significant
difference in adsorption from 60 A versus 140 A.
0 pH=4 *pH= 8
Q 20 ---------------------------------
0 50 100 150 200 250 300 3E
Pore Size (A)
Figure 4-16. Destruction of RR (10 mg/L) after 2 hours UV exposure (0.45 mW/cm2)
versus pore size.
30 60 140 320
Pore Size (A)
Figure 4-17. Adsorption of CV on gels (6% TiO2) versus pore size.
4.3 Column Studies
Since the optimal pore size for adsorption is contrary to the optimal pore size for
destruction, it was decided to test two pore sizes in a flow-through regenerative column.
These studies utilized 12% optimal titania loaded pellets of 30 A and 140 A pore sizes. A
larger pore size was not analyzed for these studies due to the frailty of the gel at pore
sizes larger than 200 A. The duplicated column exhaustion curves for each pore size are
shown in Figures 4-18 and 4-19. The duplicated runs were done with pellets made from a
Table 4-2 summarizes the exhaustion results and shows a longer time to
breakthrough of the initial concentration for the column of 30 A gels. Yet, after the 30 A
pellets reached exhaustion and the system was regenerated by photocatalysis, its level of
adsorption ability decreased after each cycle. This was also visually observed. After one
hour of UV exposure the 30 A pellets were still slightly purple, while the 140 A
pellets had regenerated and returned to its initial white state.
-- Run 1 ---*--- Run 2
0 $- ------__, -.----
0 50 100 150 200 250 300 350 400 451
Figure 4-18. Column exhaustion curve for 140 A pellet (12% TiO2).
--- --- Run1
0 100 200 300 400
Figure 4-19. Column exhaustion curve for 30 A pellets (12% TiO2).
Table 4-2. Exhaustion (C/Co = 1) times (minutes) for column studies.
Pellet Pore Size 140 A 30 A
Run 1 Exhaustion Times 36 157
Run 2 Exhaustion Times 28 132
Average Exhaustion Time 40.6 104.5
Both sets of pellets were examined during a column run after exhaustion of the
column and it was observed that the 30 A pellets were still white on the inside, revealing
that the total surface available for adsorption was not utilized. The 140 A pellets were
purple throughout their interior, revealing that the available space for adsorption had been
exhausted. Thus, the small pore size of the 30 A gels inhibits the transit of CV solution
through the porous matrix. Therefore, the 140 A sized gels exhibit the greatest efficiency
in utilizing its surface area for adsorption and its photocatalytic ability to regenerate.
Thus, the 140 A pore size was chosen as the optimal pore size for the regenerative
In order to determine if a shorter regeneration time was achievable while still
keeping the same trend as previously observed, the effect of regeneration time with UV
light was examined for the 140 A column. The time was decreased from 1 hour to 10
minutes and Figure 4-20 shows the comparison of this breakthrough curve with the 1
hour curves. The average breakthrough time was 36 minutes for the 10 minute run versus
46 for the 60 minutes run. Yet, the same trend seems to be exhibited by all 140 A
breakthrough curves. This increase in breakthrough time with each cycle shows the
benefit of using the 140 A pellets, for this trend was not observed with the 30A pellets.
--Run1 ...*... Run2 10 minute cycle
0 ,-- --------" > -,-v -* -.. -.-. -
0 50 100 150 200 250 30
Figure 4-20. Effect of regeneration time on 140 A column runs.
SUMMARY AND CONCLUSIONS
Silica-titania composites were made using a sol-gel method that allows the doping
of titania during gelation. The gelation rate was increased by use of acid catalysts and
allowed the composites to be made into pellets with sufficient dispersion of TiO2 within
the silica matrix. Varying the concentration of hydrofluoric acid allowed the
manipulation of pore size.
Three variables (TiO2 loading, curing temperature, and pore size) were
investigated to determine the best silica-titania composites for a packed column system.
These variables were optimized based upon destruction and adsorption studies and the
gels were characterized by surface area and XRD analysis. Reactive red dye (RR) was
chosen for its non-photolytic behavior in the analysis of photocatalytic ability for the
destruction studies and crystal violet dye was chosen because of its cationic nature for the
An optimal loading of 12% TiO2 was chosen based upon the results of studies
involving the destruction of RR and adsorption of CV. At higher loadings, it was
observed that the titania is agglomerating during mixing, limiting the effective surface
area available for reactions on its surface. Also, there was no trend with BET surface area
on a volume basis versus titania loading. In addition, the adsorption studies revealed a
plateau in uptake beginning at 12% TiO2.
A curing temperature of 800C showed the largest effects (a decrease of over 80%
destruction) on photocatalytic ability. At this temperature, the surface area decreased and
the phase change of anatase to rutile is evident. This resulted in a dramatic decrease in
destruction of RR. Yet, this temperature showed the highest level of adsorption. It was
concluded that the clean silica surface and more hydrophobic behavior the silica surface
displays at this temperature increased adsorption. This increase in adsorption would
decrease over time due to rehydroxylation. The 180C gel performed the best for
destruction and was chosen as the optimal curing temperature.
Concerning pore size, the highest destruction rates were obtained with large pore
sizes (> 140 A). The opposite was true for adsorption, which was greater with small pore
sizes (30 A). Column studies were done in order to determine the best pore size for a
regenerative system. The smaller pore size had the longer time to exhaustion, but did not
efficiently regenerate. The small pores inhibited the ability of water or the contaminant to
move through the matrix, i.e. the path of least resistance was around the pellet versus
through the pellet. On the other hand, the 140 A gels achieved complete regeneration
after one hour of UV exposure and efficiently utilized their surface area for adsorption. In
conclusion, the optimal pore size chosen for the regenerative flow-through column was
COMMON PREPARATION METHODS OF MIXED AND SUPPORTED OXIDES AS
DISCUSSED BY GAO AND WACHS (1999).
A.1 Mixed Oxides
A.1.1 Sol-Gel Hydrolysis
Sol-gel hydrolysis involves the acid hydrolysis and condensation of chemically
mixed Ti- and Si- alkoxides. In summary, a Ti atom acts as a substitute for a Si atom in
the silica network. It is the most widely used method due to its capability in controlling
the textural and surface properties of the mixed oxides, but the synthesis conditions can
greatly affect the homogeneity of the final product. There is a recently developed two-
stage hydrolysis procedure that solves this problem and results in the highest
homogeneity. Yet, these mixed oxides can only be obtained at low TiO2 content, less than
15 wt%. At higher Ti content, due to the larger size of the Ti atom versus Si, atom, there
is distortion of the SiO2 network. This increases the OH content on the silica surface and
affects the crystal phase of the TiO2.
A.1.2 Coprecipitation (Not Used In This Study)
Coprecipitation involves the simultaneous precipitation of Si02 and TiO2 (at high
pH) that also includes the formation of linkages with each other. This method also has
similar problems to sol-gel hydrolysis.
A.2 Supported Oxides
A.2.1 Impregnation (Not Used In This Study)
Impregnation involves the coating of silica with Ti by mixing with a titanium
precursor, such as TiC14 or a Ti-alkoxide. The following equation summarizes the
Si OH + Ti+4 Si-O-Ti
The homogeneity of the supported oxide is dependent on the concentration of surface
silanols, pretreatment temperature, and molecular size of precursor. A maximum
dispersion of TiO2 was found to be roughly 4.0 Ti atom/nm2. Consistent creation and
determination of the titania phase for catalytic purposes is a major problem with this
A.2.2 Chemical Vapor Deposition (not used in this study)
The TiO2 is deposited onto the porous silica surface from a gaseous metal
precursor. Consistent coating is difficult to attain from batch to batch and another
problem is the reduction in surface area after coating.
B.1 Methylene Blue
Empirical Formula: C16H18ClN3S
(H3C)2N Jc Sxr% N(CH3)2
B.2 Malachite Green
Empirical Formula: C23H25N2C1
B.3 Crystal Violet
Empirical Formula: C25H30N3C1
(C H3)2 N
B.4 Reactive Red
Empirical Formula: C19H12C12N607S2
Cl N CI
TEM AND SEM PICTURES
TEM of 4% TiO2 gel at 20,000x
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Danielle Julia Eller Londeree was born in June of 1979 in Deerfield Beach,
Florida. Her dad graduated from UF with an agricultural engineering degree and took
over the family business of manufacturing large-scale water pumps. Her brothers also
graduated from UF and followed in his footsteps, both receiving engineering degrees and
working for the family business. Their encouragement is what made Danielle want to be
an engineer, and following high school graduation from Zion Lutheran Christian School
in 1997, she pursued an engineering degree from UF. After getting married to Donald
Londeree in August of 2001, she received her B.S. in Environmental Engineering in
December of 2001. Then her thirst for knowledge and her desire to contribute to the
family business led her to pursue graduate school, focusing on water quality. In the
future, she hopes to use the knowledge that she has obtained in graduate school to
increase the scope of the family business and assist in solving the world's water
problems, on both a quantity and quality level.