KINETICS OF THE DECOMPOSITION AND AGGREGATION OF
SILVER NANOPARTICLES INDUCED BY HALIDE IONS
_______________
A Thesis
Presented to the
Faculty of
San Diego State University
_______________
In Partial Fulfillment
of the Requirements for the Degree
Master of Science
in
Chemistry
_______________
by
Maria Gabriela Espinoza
Summer 2012
iii
Copyright © 2012
by
Maria Gabriela Espinoza
All Rights Reserved
iv
DEDICATION
I want to dedicate this work to my parents, Luis and Dulce Espinoza, for their
ongoing support throughout my years as an undergraduate and graduate student. They have
made numerous sacrifices to help me with my education and I am very proud to have them as
my parents. I also want to dedicate this to my brother and sister, Rodolfo and Alejandra, for
being patient with me and always being there for me. They are truly my best friends. Lastly, I
could not have done it without the love and encouragement from Craig.
v
ABSTRACT OF THE THESIS
Kinetics of the Decomposition and Aggregation of Silver
Nanoparticles Induced by Halide Ions
by
Maria Gabriela Espinoza
Master of Science in Chemistry
San Diego State University, 2012
Oxidative decomposition is observed at lower halide concentrations where the rate
order of the reaction is 2/3. The more thermodynamically favorable silver halide has the
highest rate of decay: I- > Br- > Cl- > F-. The rate of nanoparticle aggregation increases
substantially when the ionic strength of the solution increases. This is due to an increase in
the screening of the electrostatic force between the particles. There is considerable interest in
nanoparticle aggregation because it intensifies the Surface-Enhanced Raman Spectrum of
species absorbed on the nanoparticle surfaces. Chloride ion is typically used for this purpose,
but, according to the Derjaguin, Landau, Verwey and Overbeek theory, the aggregation rate
of a particular type and size of nanoparticle depends only on the ionic strength, so any salt
should have the same effect. Yet, we find that chloride ion and fluoride ion have very
different critical coagulation concentrations– the concentration at which the aggregation rate
increases to the point of being diffusion-limited. We have measured aggregation rate
constants as a function of halide concentration and compared them with those calculated
using a pair-potential which combines a van der Waals attractive potential with an
electrostatic repulsive potential. This comparison shows that the different effects of the two
halide ions are likely due to their effect on the nanoparticle surface charge. The decay
reactions were also studied under various temperatures to observe any changes in the rate
decay at the different temperatures. Details of the experiments and analysis which lead us to
this conclusion will be presented.
vi
TABLE OF CONTENTS
PAGE
ABSTRACT...............................................................................................................................v
LIST OF TABLES ................................................................................................................. viii
LIST OF FIGURES ................................................................................................................. ix
ACKNOWLEDGEMENTS ..................................................................................................... xi
CHAPTER
1
INTRODUCTION .........................................................................................................1
1.1 Effect of Aggregation on Optical Properties .....................................................2
1.2 Surface Effects ...................................................................................................3
1.2.1 Solar Cells .................................................................................................3
1.2.2 Antimicrobial Properties of Silver Nanoparticles .....................................4
1.3 Surface Plasmon Resonance ..............................................................................4
1.4 Purpose...............................................................................................................6
2
METHODS ....................................................................................................................7
2.1 Materials ............................................................................................................9
2.2 Synthesis of Silver Nanoparticles ......................................................................9
2.3 Instrumentation and Kinetics Measurements ...................................................10
3
THE ADDITIONS OF SODIUM IODIDE AND SODIUM BROMIDE ...................12
3.1 Surface Plasmon Resonance Peak ...................................................................12
3.2 Oxidative Decomposition ................................................................................13
3.3 Discussion ........................................................................................................16
4
THE ADDITION OF SODIUM CHLORIDE .............................................................21
4.1 Surface Plasmon Resonance Peak ...................................................................21
4.2 Oxidative Decomposition ................................................................................22
4.3 Silver Nanoparticle Aggregation .....................................................................25
4.4 DLVO Theory For Aggregation Analysis .......................................................26
4.5 Discussion ........................................................................................................32
5
THE ADDITION OF SODIUM FLUORIDE..............................................................33
vii
5.1 Surface Plasmon Resonance peak ....................................................................33
5.2 Oxidative Decomposition ................................................................................34
5.3 Silver Nanoparticle Aggregation .....................................................................34
5.4 Discussion ........................................................................................................35
6
THE ADDITION OF SODIUM CHLORIDE FOLLOWED BY SODIUM
FLUORIDE ..................................................................................................................37
6.1 Experimental ....................................................................................................37
6.2 Results and Discussion ....................................................................................38
7
TEMPERATURE STUDIES .......................................................................................40
7.1 Varying Temperatures of the Silver Nanoparticles and Halide
Solutions ................................................................................................................40
7.2 Possible Mechanisms .......................................................................................45
7.3 Heating Silver Nanoparticles Before Initiating Decay Reactions....................47
8
CONCLUSIONS..........................................................................................................50
REFERENCES ........................................................................................................................53
viii
LIST OF TABLES
PAGE
Table 4.1. pH Changes with NaCl ...........................................................................................25
Table 5.1. Silver Halide Solubilities at 25°C ...........................................................................34
ix
LIST OF FIGURES
PAGE
Figure 1.1. Surface plasmon resonance peak of silver nanoparticles. .......................................2
Figure 1.2. Surface plasmon resonance. ....................................................................................5
Figure 2.1. Increase in decay rate due to catalysis by the light source (35mM NaBr). ...........11
Figure 3.1. Changes in the surface plasmon resonance peak after NaI addition. ....................13
Figure 3.2. Changes in the surface plasmon resonance peak after NaBr addition...................14
Figure 3.3. Log t ½ vs. sodium halide concentration for NaBr and NaI. ..................................15
Figure 3.4. Oxidative decomposition of 0.0848 mM AgNO3 and sodium citrate with
25 mM NaBr A) Slope = 1.083 x 10-2, R2 = .958 B) Slope = -1.252 x 10-3, R2
= .999. ..........................................................................................................................17
Figure 3.5. The oxidative decomposition with 25 mM NaI. A) Slope = 9.748 x 10-2,
R2 = .808. B) Slope = -6.982 x 10-3, R2 = .994. ...........................................................18
Figure 3.6. Proposed mechanism. ............................................................................................19
Figure 4.1. Changes in the surface plasmon resonance peak after 10 mM NaCl
addition to 0.113 mM silver nanoparticles. .................................................................22
Figure 4.2. The oxidative decomposition of silver nanoparticles with 1.5 mM NaCl.
A) Slope = 1.245 x 10-5, R2 = .977. B) Slope = 1.1297 x 10-4, R2 = .999....................23
Figure 4.3. Passivation with increasing [NaCl]. ......................................................................24
Figure 4.4. Aggregation of silver nanoparticles due to the addition of 45mM NaCl. A)
With a Slope of 0.216, R2 = 0.998. B) With Slope = 1.58 x 10-3, R2= 0.966. .............27
Figure 4.5. Rate order as a function of NaCl concentration. ...................................................28
Figure 4.6. DLVO potential for various concentrations of NaCl (r = 12 nm is the
smallest distance possible for particles of 6 nm radius). .............................................29
Figure 4.7. Experimental and theoretical onset of aggregation for NaCl. ...............................31
Figure 5.1. Changes in the surface plasmon resonance after NaF addition (the final
concentration of NaF is 10 mM). .................................................................................33
Figure 5.2. Experimental and theoretical onset of aggregation for NaF. .................................35
Figure 5.3. Replacement of citrate by Cl- on the surface of the silver nanoparticles. .............36
Figure 6.1. Rate constants for the reaction with 10 mM NaCl and NaF..................................38
Figure 7.1. Varying temperature effect on the half-lives of the silver nanoparticles. .............41
Figure 7.2. Aggregation at 40 mM NaCl at 8°C. .....................................................................42
x
Figure 7.3. Heated silver nanoparticle solution at 65°C. A) Increase in the surface
plasmon resonance peak after time, and B) the change in the SPR peak area. ............43
Figure 7.4. A) Heated silver nanoparticle solution at 40°C after 4 days. B) Cooled
Silver Nanoparticles at 8°C after 115 minutes.............................................................44
Figure 7.5. Ln k vs. 1/T, results in an activation of -6.561 x 104 J/mole. ................................45
Figure 7.6. Potential diagram with a negative activation energy. ............................................46
Figure 7.7. Addition of NaF following the heating of the silver nanoparticles. ......................48
xi
ACKNOWLEDGEMENTS
This thesis would not have been possible without the help from Dr. David Pullman,
Dr. Karen Peterson and group members. I want to thank Dr. Pullman for the opportunity he
gave me to be one of the first students to work on the silver nanoparticle research and
allowing me to have my own project. It is extremely exciting to see how the research has
evolved in the last five years and it has been a privilege to be working under a great mentor.
I owe most gratitude to Dr. Peterson. She has guided me through this project and I
could not have done it without her help, patience and advice. I am truly grateful to have had
the opportunity to work and learn from her.
Also, I would like to thank Mallory Hinks for all her support and encouragement
throughout the last five years. She has not only been a lab mate, but has also become a very
close friend. I want to wish her all the luck in the world as she begins her journey at U.C.
Irvine for her Ph.D. In addition, I wanted to thank Alexandra Mendoza for her help in
carrying out reactions needed for this work. She has been an incredible addition to the lab
and I am extremely fortunate to have had the chance to work with her. I wish her luck as she
pursues her doctorate degree at U.C. Los Angeles. I want to thank my family for their
ongoing faith and support and Craig for all his encouragement and positivity he has given
me.
Finally, I’d like to thank PURE Bioscience for funding the research on silver
nanoparticles.
1
CHAPTER 1
INTRODUCTION
Although the special properties of materials containing nanoparticles have been
recognized since ancient times1, their underlying structure could not begin to be understood
until the 19th century when the molecular theory of matter was becoming established. In the
1850’s, Michael Faraday rediscovered nanoparticles2, specifically synthesizing gold
nanoparticles, which absorb green light very strongly. His typical method of preparation used
phosphorous carbon disulfide to reduce a solution of gold chloride, resulting in a ruby red
solution indicating the formation of gold nanoparticles. Although, he was not able to observe
the actual size of the nanoparticle due to the lack of technological advances such as electron
microscopy, he was still the first to investigate their unique optical properties. He found that
the interaction of light differed with variations in size and shape, stating that “The mere
variation in the size of [gold] particles gave rise to a variety of resultant colours”2. He
concluded this by observing the solutions and how they look when illuminated by sunlight.
When a ray of light hit the solutions that contained finer particles, more light was transmitted
and absorbed and less light was reflected than if the solution contained nanoparticles of a
bigger size2. Faraday’s thorough research of ruby gold nanoparticle marked the birth of
modern colloidal chemistry.
The unique optical properties of metal nanoparticles are now known to be caused by
the surface plasmon resonance (SPR), whose characteristics depends on the size of the
particle and the type of metal. This resonance results in extremely high extinction
coefficients which can be exploited in many applications3. Silver nanoparticles are especially
useful because their resonance absorption occurs as a relatively sharp peak near 400 nm, a
convenient region of the spectrum (see Figure 1.1). In addition, they have important
antimicrobial properties.
Halide salts are known to affect the properties of silver nanoparticles. For example,
they are used to induce aggregation by increasing the ionic strength of solutions. The
following section explores the use of this property. They also may affect the nanoparticle
2
Figure 1.1. Surface plasmon resonance peak of silver nanoparticles.
surface directly, and this may determine the long term stability of the nanoparticle in certain
environments. Therefore, it is important to understand how the silver nanoparticles interact
when exposed to substances. Halide salts are particularly important because of their
ubiquitous presence in the environment and because silver halides are very insoluble (except
silver fluoride).
1.1 EFFECT OF AGGREGATION ON OPTICAL PROPERTIES
Silver nanoparticles have been used in Surface Enhanced Raman Spectroscopy
(SERS), which is a sensitive technique that results in the enhancement of Raman scattering
of molecules when they are adsorbed on rough, nano-sized metallic surfaces4. Roughened
metal surfaces produce SERS signals, but more recently nanoparticles have been used to
enhance the signal. The enhancement factor when using silver nanoparticles can range from
104-106, making it sensitive enough to detect single molecules.4
In SERS, the signal is enhanced by adding nanoparticles to the substance being
analyzed. Aggregating the silver nanoparticles has shown to increase the signal. The
aggregation of the silver nanoparticles is primarily done by introducing solutions of sodium
halides. The addition of the sodium halides lead to changes on the surface of the silver
nanoparticles, therefore resulting in an increase of signal and greater enhancement.
Another application of silver nanoparticles, which exploits their optical properties, is
their use as DNA biosensors. A study at Drexel University consisted of injecting silver
3
nanoparticles into targeted cancerous cells or different molecules. Then a halide solution was
introduced to aggregate the silver nanoparticles. This resulted in a change in color for each of
the nucleotides. The four DNA nucleotides exhibited different colors when the silver
nanoparticles are attached, enabling the detection of the DNA nucleotides5.
1.2 SURFACE EFFECTS
The surface structure of nanoparticles and the effect of the environment on that
structure are likely to be critical in their use in optical applications and in medicine, but there
is very little research in this area. Not only can adsorbed species affect their applicability
directly, but they can also affect their long term stability due to either aggregation or
oxidative decomposition.
1.2.1 Solar Cells
The SPR of silver nanoparticles makes them ideal in the enhancement of solar cells.
Like plants, solar cells turn light into energy. Plants do this inside vegetable matter, while
solar cells do it in a semiconductor crystal doped with extra atoms. Current solar cells cannot
convert all the incoming light into usable energy because some of the light can escape back
out of the cell into the air. Additionally, sunlight comes in a variety of colors and the cell
might be more efficient at converting bluish light while being less efficient at converting
reddish light.
The silver nanoparticle approach in enhancing solar cells may serve as a viable
solution to this problem. The electron oscillations on the surface of the nanoparticle are
excited when incoming light hits the surface of the nanoparticle, and at a resonant
wavelength, the scattering of light is very strong, therefore, enhancing the efficiency of the
solar cell. To make use of silver nanoparticles, they must interact with the rest of the cell.
Over coating the metal particles on a substrate can change the resonance and can also affect
the resonance frequency. Red shifting of resonance leads to an increase in scattering cross
section at longer wavelengths and an increase of the absorption of solar cells because of the
incident band gap of silicon. The silver ions, Ag+, are reduced on the silicon wafer surface by
injecting into the valence band of silicon. Silver has a higher electronegativity than silicon,
therefore, strongly attracts electrons from the silicon to become negatively charged6. The
4
nanoparticle approach has been growing in popularity because silver nanoparticles can
increase the efficiency of the solar cells up to 30 percent.
1.2.2 Antimicrobial Properties of Silver Nanoparticles
Silver nanoparticles have unique antibacterial properties that make them ideal
candidates for medical applications. The antibacterial properties of silver have been known
for thousands of years beginning with Ancient Greece. The Greeks coated their pots and
utensils with silver and determined that their mouths were kept healthier. Silver has long
been known for its toxicity against a wide range of micro-organisms; and for this reason
silver nanoparticles have been widely used in various antibacterial applications. It has been
proposed that ionic silver strongly interacts with the thiol groups of vital enzymes and
inactivates them. Silver nanoparticles act as catalysts, disabling the enzyme that bacteria,
viruses and fungi need for their oxygen metabolism7. Silver nanoparticles have been used
medically to treat burns and a variety of infections. In a recent study, it was even shown that
embedding tissue grafts with silver nanoparticles was more effective than the widely used
silver sulfadiazine salt. They were able to show that besides exhibiting strong antibacterial
activity, silver nanoparticles also showed signs of decreasing inflammation7. Understanding
the chemical properties and kinetic behavior of silver nanoparticles is critical and extremely
important for their possible medical applications.
1.3 SURFACE PLASMON RESONANCE
Approximately fifty years following Faraday’s gold nanoparticle work, Gustav Mie
presented a theoretical explanation for the color of Faraday’s gold nanoparticles. In 1908
Gustav Mie8 presented solutions to the Maxwell equations to describe the extinction
spectrum, which is the sum of the amount of light scattered and absorbed, of spherical
nanoparticles of arbitrary size. For particles much larger or much smaller than the
wavelength of the scattered light, there are simple and excellent approximations that suffice
to describe the behavior of the system. For example, discrete dipole approximations (DDA)
are used to solve the equations for particles of other geometries such as cubes and rods3.
These calculations clearly show that the amount of light being absorbed and scattered is
dependent on the shape and size of the silver nanoparticles.
5
What is responsible for the unique interaction of light with metallic nanoparticles of
different sizes and shapes? When a small spherical nanoparticle is irradiated with resonant
incident light, the small electron clouds on the surface of the nanoparticles begin to oscillate
coherently (see Figure 1.2). The oscillation frequency, which is dependent on electron
density, the effective electron mass, and the shape and size of the charge distribution9, is
called the surface plasmon resonance frequency. The resonance frequency can change when
species are adsorbed on the surface of the metal particle. When incident light strikes the
surface, the localized surface plasmon resonance is excited3. Thus, silver nanoparticles of
different sizes and shapes scatter and absorb light differently and exhibit different colors, as a
result of changes in the surface plasmon resonance.
Figure 1.2. Surface plasmon resonance.
The area of the surface plasmon resonance peak is proportional to the number of
electrons. This is a useful property for monitoring the total amount of silver atoms contained
in the nanoparticle solutions. Therefore, by measuring the decay of the SPR peak, the decay
rate of the nanoparticles can be determined. This decay can be due to the nanoparticles
becoming smaller or to their loss by an aggregation process. The shape of the SPR peak is
strongly dependent on the immediate environment of the nanoparticle. Damping of the
surface plasmon peak can be observed in the presence of different surface species. For
example, Mulvaney studied the damping effects due to the presence of iodide ions which
show a red shift in the surface plasmon resonance peak10. By monitoring the surface plasmon
resonance peak, changes of the surface of the silver nanoparticles can be probed.
6
1.4 PURPOSE
With the increasing usage of silver nanoparticles in commercial products, it is
important to study how they react in the environment. Therefore, it is important to understand
how the silver nanoparticles react under various environmental conditions. The main focus of
this work was to quantitatively investigate the physical properties of silver nanoparticles in
different halide solutions, particularly with respect to various kinetic processes when reacted
with NaI, NaBr, NaCl and NaF. We observed how long the silver nanoparticles survived
when subjected to a range of physical and chemical conditions, such as various
concentrations of the halide solutions and varying temperatures.
When the silver nanoparticles are reacted with the sodium halides, two primary
kinetic processes were observed; 2/3rd order oxidative decomposition and 2nd order
aggregation. Oxidative decomposition typically occurs at lower concentrations of sodium
halides. Decomposition of the nanoparticles is evident based upon the rate order, a narrow
width of the surface plasmon resonance peak, and the resultant reddish precipitate.
Aggregation occurs at higher concentrations of sodium halides. The evidence for aggregation
are the rapid decrease in the surface plasmon resonance peak, an increase in the rate constant,
2nd rate order and the presence of a gray precipitate. The kinetic processes will be explained
in further detail in the ensuing chapters.
7
CHAPTER 2
METHODS
There are several different methods for producing different sizes and shapes of silver
nanoparticles, each with its advantages and disadvantages. The morphology and size of the
silver nanoparticles have been shown to be dependent on temperature, the capping agent
used, and the concentration of the silver source. Some of these methods will be summarized
before describing in detail the method used in this work.
One of the oldest methods for synthesizing silver nanoparticles is a variation of the
simple Turkevich method, which is used to synthesize monodisperse gold nanoparticles11-12.
In this method, chloroauric acid is reduced by sodium citrate, producing spherical
nanoparticles of about 10 nm in diameter. The citrate ion serves not only as the reducing
agent but also as the capping agent of the nanoparticles, enabling them to be electrostatically
stabilized. The variation of this technique, to produce silver nanoparticles, is known as the
Lee and Meisel method13. A silver nitrate solution is brought to a boil, and the silver ions are
reduced by the dropwise addition of trisodium citrate. Although simple and straightforward,
the Lee and Meisel method produces nanoparticles with a broad range of sizes, and this is a
major disadvantage 14.
Sun and Xia synthesized silver nanoparticles using the polyol method12, which is a
method for producing nanoparticles with unique morphologies. This method involves the
reduction of silver nitrate by ethylene glycol in the presence of the capping agent poly(vinyl
pyrrolidone) (PVP) at relatively high temperatures, between 120°C and 190°C. They found
that the shape of the nanoparticles could be varied by varying the temperature at which they
were synthesized. They also studied the effect of decreasing the concentration of silver
nitrate, and found, for example, that lower concentrations of silver nitrate led to nanowires as
the dominant shape. Their experiments clearly show that the morphology of the silver
nanoparticles is influenced by the experimental conditions and method of synthesis12.
Jiang et al. studied the role of temperature in the growth of silver nanoparticles 15. In
their study, silver nanoparticles were prepared via a synergenic reducing approach using two
8
or three different reducing agents: citric acid, L-ascorbic acid and sodium borohydride at
room temperature. Prepared in this way, silver nanoplates formed were flatter in shape than
the usual spherical nanoparticles and had extremely large adsorbing and scattering crosssections across the visible and near-IR region of the spectrum. With increasing temperatures,
the triangular silver nanoplates grew from 90 nm to 180 nm, as a result of a fusion
mechanism where multiple nanoplates fuse together to form one of a bigger size.
Leopold and Lendl developed a method of the reduction of silver nitrate using
hydroxylamine hydrochloride 16. This method reduces silver nitrate with an alkaline solution
of hydroxylamine hydrochloride. The silver nitrate solution is mixed as the hydroxylamine
hydrochloride is added dropwise. By changing the mixing order, or the rate at which the
hydroxylamine hydrochloride is added, one can control the size and the dispersion of the
silver nanoparticles. By changing the mixing order and the mixing rate, different particle size
distributions were observed, with average sizes ranging from 23 to 67 nm. The nanoparticles
prepared in this manner can be used immediately for SERS. Overall, this procedure is fast,
simple and produces nanoparticles that are comparable with those prepared using the LeeMeisel method, but prepared in a much simpler manner.
Van Hyning et al. have conducted extensive studies on the Creighton method in
which a silver salt (in his case, silver perchlorate) is reduced by sodium borohydride 17. They
divided the system of reactions into three different stages based on their visual observations
of the silver nanoparticles. The reaction was initiated by adding an excess of sodium
borohydride at 15°C to the silver source; this was the first observation stage. After the first
stage, the surface plasmon peak was initially broad and then sharpened at 400 nm, which
indicated that small nanoparticles of a size of 1.5-2 nm in diameter were produced. In the
intermediate stage, the nanoparticles grew to about 5-15 nm in size. The last stage consisted
of the reaction between the borohydride and water. As all the borohydride was consumed,
aggregation commenced. This resulted in the surface plasmon peak broadening and shifting
due to the rapid increase in size. In this final stage, the solution darkened.17 This work
showed that small variations in the synthetic method can lead to large changes in the
products.
The most common method used for synthesis of silver nanoparticles, and the method
used in this work, is a variation of the Creighton method. In this variation, a capping agent is
9
used to prevent aggregation, which is a problem in the regular Creighton method 15. The
capping agent stabilizes the surface of the nanoparticle and stops the growth of the
nanoparticle before aggregation can occur. In our work, we use silver nitrate as the silver
source, with sodium citrate used as the capping agent. Sodium borohydride is still used as the
reducing agent.
2.1 MATERIALS
For the synthesis of silver nanoparticles, an analytical grade silver nitrate (99.9%)
along with the sodium halides used for the decomposition and aggregation of the silver
nanoparticles were purchased from Fisher Scientific. Trisodium citrate dehydrate (>99.5%)
and sodium borohydride (>98%) were purchased from J. T. Baker and Sigma-Aldrich, and
sodium iodide (>99%) from Acrōs Organics. All chemicals were used without further
purification. All nanoparticles were synthesized in 5 MΩ deionized water obtained in the
laboratory.
2.2 SYNTHESIS OF SILVER NANOPARTICLES
In this study, solutions of silver nanoparticles are synthesized via a variation of the
Creighton method. The silver cation is reduced by sodium borohydride in the presence of the
capping agent, sodium citrate. The sodium citrate adsorbs electrostatically on the surface of
the nanoparticles, and serves to stabilize the particles. In a 20 mL scintillation vial equipped
with a stir bar, 0.4 mL of 4.5 mM AgNO3 and 0.4 mL of 4.5 mM sodium citrate are mixed.
The solutions are diluted to a final volume of 10 mL. Then 2 mL of 0.45 mM sodium
borohydride is added dropwise to the silver nitrate and sodium citrate mixture under constant
stirring at room temperature. It is important to note that each batch of silver nanoparticles is
synthesized in the same manner to assure that the particles would be consistent. Transmission
electron microscopy indicates that the nanoparticles are spherical and have a diameter of 1015 nm. Based on this approximate size, we estimate that each nanoparticle contains about
53,000 silver atoms, indicating a nanoparticle concentration of about 2.13 x 10-9 M.
As the sodium borohydride is added, the solution changes from colorless to bright
yellow, which corresponds to the surface plasmon resonance around 394 nm. The narrow
linewidth of 50-60 nm indicates a narrow distribution in the size of the nanoparticles. The
peak absorbance is consistently between 1.0 and 1.5 AU. After the synthesis of the
10
nanoparticles in the scintillation vials, the bright yellow solutions are stored for at least four
days before being used in experiments in order for the surface to become stabilized. Because
silver nanoparticles are photosensitive, all silver nanoparticle solutions are stored in the dark
and are stable for months.
After the silver nanoparticles have aged for at least four days, sodium halide solutions
are added, bringing the total volume in each vial to 16 mL, to initiate the decay reactions.
The final concentrations are 0.113 mM for silver atoms, 0.113 mM for sodium citrate and
0.056 mM for sodium borohydride. The final halide concentrations vary from 0.10 mM to
120 mM.
2.3 INSTRUMENTATION AND KINETICS MEASUREMENTS
To monitor changes in the silver nanoparticle solutions throughout the reaction, we
used either a Jasco V-670 Spectrophotometer or a Hewlett Packard Diode Array
Spectrophotometer (HP-8452A). The integration time using the HP-8452A to experiment
was set to 0.19 s, and the spectra were collected. Due to the photosensitivity of the silver
nanoparticles, it is important to use a small integration time. The absorption spectra in the
range of 190–820 nm were collected using a quartz cuvette with water as the reference. After
the addition of the sodium halides, a decay in the surface plasmon resonance peak was
observed over time. This decay was monitored by measuring the absorbance of the SPR peak
at about 394 nm as a function of time.
During the experiments it was observed that the reaction between silver nanoparticles
and sodium halide solutions exhibited photosensitivity. Originally, the same aliquot of
sample was used to measure the decay reaction; however, we observed that the light source
of the spectrometer was catalyzing the reaction. In order to test this, we compared the rates
for the decay of the surface plasmon peak when using the same aliquot sample for each scan
and the rate when using different aliquots of sample for each scan (cleaning the cuvette
between scans. After data analysis in Excel, it was determined that the rate of decay was
greater when using the same sample for each scan. Figure 2.1 shows a comparison of the
decay rate for a 35 mM NaBr reaction with the silver nanoparticles. For the reaction using
the same aliquot of the solution, the rate was found to be an order of magnitude faster than
when fresh aliquots were used for every new absorption measurement. This result indicates
11
Figure 2.1. Increase in decay rate due to catalysis by the light source
(35mM NaBr).
that the reaction was being catalyzed by the light source of the spectrometer. Due to this
observation, our protocol was to use fresh aliquots of sample for each separate measurement.
In addition to studying the effect of various halides on the silver nanoparticles,
temperature studies for the reaction of the silver nanoparticle solution with 40 mM sodium
chloride at temperatures ranging from 4- 40°C were performed and monitored as a function
of time. The solutions were heated or cooled using an Echotherm chilling/heating dry bath
from Torrey Pines Scientific, Inc. Once they reached the specific reaction temperature, the
sodium halide and silver nanoparticles solutions were mixed together and the absorption
spectra were measured as a function of time. Similar to the decay measurements previously
discussed, the decay of the surface plasmon resonance peak was measured using a UV-Vis
spectrometer as a function of time in the wavelength range 190-820 nm.
12
CHAPTER 3
THE ADDITIONS OF SODIUM IODIDE AND
SODIUM BROMIDE
In order to measure the kinetics of the decomposition and aggregation of silver
nanoparticles, sodium halides were added to the solution. In this section, the reactions
between silver nanoparticles and sodium iodide and sodium bromide are discussed. To begin
the experiments, concentrations in the range 6-85 mM of sodium bromide or sodium iodide
were added to the silver nanoparticle solutions. The decay of the surface plasmon resonance
peak in the UV-Visible spectrum was then monitored as a function of time.
The peak absorbance, A, is proportional to the concentration of the silver
nanoparticles if the full width half maximum (FWHM) does not change significantly in time.
The presence of oxidation can be observed by the linearity of the A1/3 vs. time plot, which
indicates whether the rate order for the reaction is 2/3. In a 1/A vs. time plot the linearity
indicates a second order reaction and suggests an aggregation process in which the
combination of two particles is the first step occurring in the reaction.
3.1 SURFACE PLASMON RESONANCE PEAK
Silver nanoparticles typically have an absorption peak at a wavelength of 393 ± 2 nm.
The additions of sodium halides can result in different changes of the absorption peak
depending on the halide added. We observe not only changes in the rate order, but also
changes in the surface plasmon resonance peak which can be attributed to changes occurring
on the surface of the nanoparticle. In past research, Heinglein et al. showed that in the
presence of oxygen, the nanoparticles develop an oxide layer of Ag2O. The presence of the
oxide layer broadens the FWHM of the peak from about 25 nm to 50 nm, and shifts the peak
towards the red region from 382 nm to 394 nm.18
Immediately following the addition of sodium iodide, the surface plasmon peak
changed as seen in Figure 3.1. The peak absorption decreased while the peak width
increased. We also observed a significant red shift of the peak from 392 nm to about 420 nm.
13
Figure 3.1. Changes in the surface plasmon resonance peak after NaI
addition.
The color of the solution changes from yellow to orange because of the broadening of
the peak. Mulvaney suggested that the red shift and peak shape changes are due to the
presence of silver iodide (AgI). He concluded that an extremely thick layer of AgI would be
required to broaden the surface plasmon resonance peak, while also seeing a new peak form
at around 420 nm 10. However, we observed the formation a second peak, AgI, at around 428
nm, but not until the reaction was approximately 50% complete, where the adsorption peak
had decreased by one half, as shown in Figure 3.1.
In the reaction between silver nanoparticles and sodium bromide, the surface plasmon
peak exhibited less broadening than NaI (Figure 3.2). There is a small red shift to 402 nm, a
slight decrease in the absorption peak height, and a slight increase in the peak width. These
relatively small changes indicate that the bromide ions are modifying the surface of the silver
nanoparticle by replacing the citrate ions or oxidizing the oxide layer, but to a lesser degree
than with the addition of sodium iodide.
3.2 OXIDATIVE DECOMPOSITION
After the initial changes in its shape, the surface plasmon peak begins to decay
through time. The rate of decay is dependent on the concentration of sodium halide added to
the silver nanoparticle solutions. For NaI and NaBr, the decay rate is fast at high and even
low concentrations. Sodium iodide decay reactions have rate constants about one order of
magnitude greater than sodium bromide. A useful indication of a 2/3rd order chemical
14
Figure 3.2. Changes in the surface plasmon resonance peak after NaBr
addition.
reaction is the half-life, t1/2, which is the time it takes for its absorbance to fall to half the
initial value. From the rate constants of a 2/3rd order reaction, the half-life can be calculated
using:
t1 / 2
22 / 3 − 1
=
2
n −1
( )k[ A]o
3
The half-life for reactions that are slow can be calculated by extrapolating from the
data the time (in minutes) of where the absorbance is expected to be at half its original value.
Figure 3.3 shows the difference in the half-lives of the silver nanoparticles in reaction with
the sodium bromide and iodide solutions. The nanoparticle half-lives are longer in the case of
sodium bromide.
Additional kinetic information can be determined for these reactions by deriving the
integrated rate laws. At low concentrations of sodium bromide and at all concentrations of
sodium iodide, oxidative decomposition was shown to be the dominant kinetic behavior.
First, to determine the rate order of these decay reactions, we assumed that the rate of silver
lost is proportional to the surface area available for the reactions, as seen in Equation 3.1.
dN Ag
dt
= −keff S
(3.1)
15
Figure 3.3. Log t ½ vs. sodium halide concentration for NaBr and NaI.
Where the total surface area of the nanoparticle is S, and the number of atoms of
silver is N Ag . Also, in the presence of a reactive species in the solution, O2, the effective rate
constant, keff , remains constant when the concentration of oxygen remains constant
throughout time. The surface area, S, is constant for reactions on flat surfaces, but for small
nanoparticles, the area will decrease as the decay reaction continues with time. The surface
area will be dependent on the amount of silver moles, N Ag . Equation 3.2 is derived from the
4
volume ( πr 3 = [ Ag1 ]M /[ Ag n ]o ρN A ) and surface area ( [ S ] = 4πr 2 [ Ag n ]N A ) of the silver
3
nanoparticle.
 3[ Ag1 ]M
S = 
ρ




2/3
(4π [ Ag n ]0 N A )1/ 3
(3.2)
The surface are is dependent not only on the molecular mass, M, but also on the
initial concentrations of silver nanoparticles, [ Ag n ]0 , and the changes in concentration of
silver atoms, [ Ag1 ] . The surface area is also inversely dependent on the density of silver, ρ.
Assuming that the amount of silver is proportional to the maximum absorbance of the surface
plasmon resonance peak, the amount of silver can be monitored using UV-Visible absorption
spectroscopy. Therefore, the decay of the absorbance will follow a 2/3 order rate law given
by Equation 3.3.
16
dA p
dt
= − k eff AP2 / 3
(3.3)
Integration of Equation 3.3 gives Equation 3.4:
1
A 1P/ 3 − A 1P,/ o3 = − k eff' t = −k obs t
3
(3.4)
Plotting A1P/ 3 vs. time, t, should result in a line with a corresponding slope equal to the
negative of the rate constant, kobs . For concentrations of sodium bromide below 50 mM, the
rate followed a 2/3rd order, shown in Figure 3.4. For all concentrations of sodium iodide, the
fit was linear, indicating that the reaction upon the addition of NaI does indeed follow a 2/3
order oxidative decomposition, shown in Figure 3.5.
3.3 DISCUSSION
Silver nanoparticles synthesized using silver nitrate, sodium citrate and sodium
borohydride are very stable. However, upon the addition of sodium halides, the silver
nanoparticles begin to decompose. In this chapter, we focused on determining the kinetics of
the processes that occur with the reactions between silver nanoparticles and sodium bromide
or sodium iodide. Interesting observations are made. The surface plasmon resonance peak
changes in amplitude, width and position immediately following the sodium halide solutions.
We hypothesize that these initial changes in the surface plasmon peak can be attributed to the
modification of the surface of the silver nanoparticle. On the basis of this hypothesis, we
have proposed the mechanism shown in Figure 3.6.
The first step of the mechanism is the rapid replacement of the citrate ions and/or
oxide layer by the halide which occurs in less than one minute following the addition of
sodium halide to the silver nanoparticle solution. The next step (2) is the ejection of electrons
into solution followed by the slow oxidative decomposition of the silver nanoparticle (3a).
During the slow decomposition, AgX is lost into solution while forming a new layer of the
silver halide on the surface of the nanoparticle. For all concentrations of sodium iodide and
concentrations below 50 mM sodium bromide, the silver nanoparticles undergo slow
oxidative decomposition. There is no sign of any aggregation occurring even at higher
concentrations of sodium iodide. However, at 50 mM NaBr, there begins to be evidence of
aggregation, where the reaction order changes from 2/3 to 2. Although silver nanoparticle
17
Figure 3.4. Oxidative decomposition of 0.0848 mM AgNO3 and
sodium citrate with 25 mM NaBr (A) Slope = 1.083 x 10-2, R2 = .958
(B) Slope = -1.252 x 10-3, R2 = .999.
18
Figure 3.5. The oxidative decomposition with 25 mM NaI. (A) Slope =
9.748 x 10-2, R2 = .808. (B) Slope = -6.982 x 10-3, R2 = .994.
19
Figure 3.6. Proposed mechanism.
reactions with NaBr follow a 2/3 order oxidative decomposition up until 50 mM, the decay
rate was about one order of magnitude slower than that of NaI.
On the surface of the silver nanoparticle, there are many oxidation and reduction
reactions present. Therefore, the surface of the silver nanoparticle contains several anodic
and cathodic sites. In the presence of O2 and X-, the cathodic and anodic reactions are shown
in Equations 3.5 and 3.6, respectively. The complexation of the silver halide salt, AgI, gives
an overall standard potential for the cell of 0.5532 V, which indicates a favorable reaction.
Upon the addition of sodium bromide, the standard electrode potential for the cell is 0.3297
V for the complexation of AgBr.
1
k1
O2 + H 2O + 2e − →
2OH −
2
k2
2 Ag + 2 X − →
2 AgX ( s ) + 2e −
rate =
4k1 k 2 S [O2 ][ X − ]
4k1 [O2 ] + k 2 [ X − ] ; X = Cl-, I-, Br- or F-
(3.5)
(3.6)
(3.7)
As stated earlier, the addition of sodium halides affect the surface plasmon resonance
peak. Immediately following the addition of sodium iodide, the surface plasmon peak
broadens and decreases in absorbance. This is evidence of the iodide ion replacing the citrate
and the oxide layer on the surface of the nanoparticle. However, the addition of sodium
20
bromide has an intermediate effect compared to NaI. There is only a slight red shift, and
almost no change in the absorbance. This indicates that the surface of the silver nanoparticle
is not changing as much compared to NaI.
For all concentrations of sodium iodide, the rate of the oxidative decomposition
followed a 2/3rd order. Therefore, it can be concluded that the rate of decomposition is
dependent on the halide concentration. Sodium bromide clearly remains 2/3rd order until 50
mM, where the rate of the reaction changes to 2nd order and the aggregation process begins to
dominate. After measuring the rate of decay of the surface plasmon peak, rate constants were
calculated for NaI concentrations ranging from 0.06-45 mM. Although the addition of Icauses changes of silver nanoparticle surface, I- does not cause any aggregation of silver
particles. It is not yet clear from the literature why the bromide ion can lead to aggregation
but the most strongly binding halide, I-, does not lead to any aggregation.
21
CHAPTER 4
THE ADDITION OF SODIUM CHLORIDE
One might expect that the reactions between the silver nanoparticles and all sodium
halides to be the same because all have the same charge and same effect on ionic strength.
The main difference is their size, which affects the strengths of interactions. In addition, the
crystal lattice is increasingly stabilized in the order of: AgF, AgCl, AgBr and AgI. For these
reasons, we observe differences in their effects on silver nanoparticles. In this chapter, we
will show that the rate order, onset of aggregation, and changes in the surface plasmon
resonance differ significantly in the case of sodium chloride relative to bromide and iodide.
4.1 SURFACE PLASMON RESONANCE PEAK
Futamata and Maruyama studied the effects of halides on the surface enhanced
Raman spectrum of rhodamine 6G 19. They found that a layer of silver chloride, AgCl, on the
bare surface of silver nanoparticles led to the surface plasmon peak shifting to the red, but
not to the same extent as for a silver oxide layer, Ag2O. They attributed this to the
translucence of AgCl in the visible spectrum compared to the small absorbance of Ag2O.
Since we make the silver nanoparticles in the presence of air, and therefore O2, the surface is
at least fractionally covered by an oxide layer. In agreement with the research done by
Futamata and Maruyama, when the oxide is replaced by chloride, the red shifting would be
less and we should actually observe a blue shift.19
Figure 4.1 illustrates that the addition of 10 mM sodium chloride to 0.113 mM
solution of silver nanoparticles increases the peak absorbance of the surface plasmon band,
decreases the width of the peak, and blue shifts the peak by 1-2 nm. This indicates that Cl- is
on the surface of the nanoparticle and is replacing the citrate and oxide layer. The significant
changes of the surface occur within a minute after the addition of sodium chloride, although
slight changes continue to occur for hours following the initiation of the decay reaction.
As seen by the modification of the SPR peak, it is evident that the surfaces of the
silver nanoparticles are changed under various environments. However, the area of the curve
22
Figure 4.1. Changes in the surface plasmon resonance peak after 10 mM
NaCl addition to 0.113 mM silver nanoparticles.
should only be dependent on the number of free electrons in the particle. Assuming that each
silver atom in each nanoparticle contributes one free electron, the area of the curve will be
proportional to the number of atoms in each nanoparticle 20. A 5% decrease in the area is
observed after the addition of NaCl, indicating that there are some silver atoms on the surface
of the nanoparticles which are being quickly oxidized. Thus it appears that, although we let
the nanoparticle solutions age for at least four days, the surface of the silver nanoparticle is
not completely oxidized and addition of NaCl increases the rate of oxidation.
4.2 OXIDATIVE DECOMPOSITION
As for NaI and NaBr, oxidative decomposition is observed for NaCl, but the rate is
much slower. After the initial changes in the spectrum, the SPR decreases in time at a rate
which initially depends on the chloride concentration. It then reaches a maximum around 1.0
mM, after which the rate decreases to a very low value, close to that of the unperturbed
nanoparticle. Using the kinetic analysis described in Chapter 3, we determined the order of
the reactions between the silver nanoparticles and NaCl. A1/3 vs. time is plotted in Figure 4.2
and is linear as expected for a reaction of 2/3rd order. Also, a plot of 1/A vs. time is given to
show that it is not linear; therefore, the reaction doesn’t proceed as 2nd order. Above 2.0 mM
NaCl the reaction becomes very slow and the order was not determined. Figure 4.3 shows the
rate constant as a function of NaCl concentration in this region. Between 6 and 25 mM (not
shown) the rate constant remains slow. The 2/3rd order is observed up to 25 mM, where
23
Figure 4.2. The oxidative decomposition of silver nanoparticles with
1.5 mM NaCl. (A) Slope = 1.245 x 10-5, R2 = .977. (B) Slope = 1.1297 x
10-4, R2 = .999.
24
Figure 4.3. Passivation with increasing [NaCl].
suddenly the rate order changes to 2nd order. The analagous NaBr reactions show similar
behavior (see section 4.3). The change in the rate of the reaction is observed with the
presence of aggregation and an increase in the rate constant at higher concentrations of
sodium chloride.
The complexation of silver chloride causes the reaction to be favorable with a cell
potential of 0.1787 V:
1
k1
2OH −
O2 + H 2O + 2e − →
2
k2
2 Ag + 2Cl − →
2 AgCl ( s ) + 2e −
If we first ignore the insolubility of the silver salt, we can assume that the chemical
reaction at the anode is the rate limiting reaction. This suggests that the surface has become
passivated, hindering further decomposition. At concentrations up to 3 mM, the oxidative
decomposition rate is at its highest and then starts to decrease due to passivation. Passivation
occurs on the surface on the silver nanoparticle when an insoluble product, in this case AgCl,
builds up and inhibits the reactants from reaching the surface of the silver nanoparticle,
lowering the expected rate constant. If passivation is not occurring, the rate constants should
level off as shown in Figure 4.3. In addition, passivation was not observed in the reactions
between the silver nanoparticles and sodium bromide and sodium iodide. It is interesting that
AgCl shows evidence of surface passivation even though AgI is more insoluble. One possible
25
explanation may be the size of I- (206 pm) compared to Cl- (167 pm). Iodide anion, being
larger, may not fit as well between the atomic spacing of the silver lattice compared to a
smaller chloride ion.
Further evidence for the decomposition pathway is given by monitoring the
production of OH- in time. This experiment was done for the reaction of silver nanoparticles
with various concentrations of NaCl and NaI. Table 4.1 shows the changes in the pH of the
silver nanoparticles solutions with sodium chloride. At a NaCl concentration of 5 mM, the
pH increased from 6.14 to 9.13 after 13 days, indicating the production of OH-. This is
expected if Equations 3.5 and 3.6 apply. At a NaCl concentration of 45 mM, the pH did not
change significantly; therefore, we can conclude that there is essentially no evidence of
oxidative decomposition at this higher concentration. In fact, we show in the next section that
aggregation becomes the dominant mechanism at higher concentrations.
Table 4.1. pH Changes with NaCl
5 mM NaCl, Time (hrs.)
pH
45 mM NaCl, Time (hrs.)
pH
0
6.14
0
7.18
1.45
6.22
0.03
7.23
20.32
6.43
0.13
7.22
118.42
7.48
0.22
7.14
120.13
8.03
0.30
7.18
140.12
8.40
0.42
7.15
184.63
8.81
1.05
7.27
258.03
9.03
1.18
7.20
308.55
9.13
20.35
7.36
4.3 SILVER NANOPARTICLE AGGREGATION
As sodium chloride concentrations increase above 25 mM, there is a sudden increase
in the rate of decay. In addition, we find that the reaction becomes 2nd order with respect to
the particle concentration. The rate law is given by the following equation:
d [ Ag n ]
= k obs , 2 [ Ag n ] 2
dt
(4.1)
26
If the silver nanoparticle size is uniform, the concentration of silver nanoparticles,
[ Ag n ] , is proportional to the concentration of silver atoms and this, in turn, is proportional to
the SPR peak. Thus, Equation 4.2 applies.
dA p
dt
= k obs , 2 A p
2
(4.2)
Integrating Equation 4.2:
1
1
=
+ kt
[ Ap ] [ Ap ]0
Thus a plot of 1/A vs. time should be linear if the reaction is second order with
respect to Ap.
Figure 4.4 shows the plot of 2nd order aggregation for 40 mM NaCl. Comparing this
to a plot of A1/3 vs. time, it is clear that the former gives a much better fit. This rapid increase
in the rate constant and change in rate order for the reaction indicate the onset of aggregation.
For NaCl, the onset of aggregation is observed at 25 mM (Figure 4.5). The 2nd order rate
suggests that the initial rate limiting step is the association of two particles.
4.4 DLVO THEORY FOR AGGREGATION ANALYSIS
The analysis of the particle aggregation observed in our studies can be interpreted in
terms of the Derjaguin, Landau, Verwey and Overbeek (DLVO) theory for particle stability.
The theory describes the force between charged surfaces interacting with one another. It
combines the effects of the van der Waals attraction, VvdW, and Coulomb repulsion, Vcoul.
The overall potential energy of interaction between the two charged silver nanoparticles is
the sum of both the attractive van der Waals component and the Coulomb (electrostatic)
repulsive component:21
U (r ) = VvdW (r ) + VCoul (r )
(4.3)
The van der Waals interaction potential is dependent on two quantities: the size of the
silver nanoparticle with radii, a , and the distance between the two particles’ centers, r.
VvdW (r ) = −
AH
6
 2a 2

2a 2
4a 2

+
+
−
ln
1
 2
2

r2
r2

 r − 4a



(4.4)
27
Figure 4.4. Aggregation of silver nanoparticles due to the addition of
45mM NaCl. A) With a Slope of 0.216, R2 = 0.998. B) With Slope =
1.58 x 10-3, R2= 0.966.
28
Figure 4.5. Rate order as a function of NaCl concentration.
AH is the Hamaker constant; it is a type of van der Waals force constant which
depends on the metal in the particles. The Hamaker constant for silver is calculated to be
around 3x10-19 J, but may be an order of magnitude lower if molecules are adsorbed on the
surface of the silver.22
The Coulombic repulsion component is dependent on both the charge of the silver
nanoparticle and the inverse of the Debye-Hückel screening length, κ, as shown in Equation
4.5.
2
e 2 Z 2  eκa  e −κr
VCoul (r ) =


4πε 1 + κa  r
(4.5)
The Debye-Hückel screening length is also dependent of the ionic strength of the
solution, I:
κ2 =
4e 2 I 2
e2
=
εk B T εk B T
N
∑n
j =1
j
z 2j
(4.6)
If the salt concentration reacting in the solution is increased, the Debye-Hückel
screening length decreases to a point where the electrostatic force is completely screened and
aggregation becomes diffusion-limited.21 The concentration at which this occurs is called the
critical coagulation concentration. If the potential barrier is reduced to less than the thermal
energy, kBT (Boltzmann constant and absolute temperature), every collision will cause the
nanoparticles to join and lead to fast aggregation. However, if the potential barrier is greater
29
than the thermal energy, many collisions are required before two nanoparticles can come
together and the system experiences slower aggregation.
It is important to understand how the potential barrier of the silver nanoparticles may
be affected by changes in the concentration, surface charge, the distance between the radii
centers, and the Hamaker constant. Figure 4.6 shows the potential barrier of the reactions
with different concentrations of NaCl. As the concentration increases, the potential barrier
decreases which leads to faster aggregation, which we experimentally observe. Also, when
the surface charge of the nanoparticle decreases and the Hamaker constant and the
concentrations remain constant, the potential barrier also decreases and leads to faster
aggregation. The opposite is true, when the charge increases and the potential barrier
increases, which result in slower aggregation kinetics.
20
30mM NaCl
35 mM NaCl
15
40 mM NaCl
UDLVO/kBT
10
5
0
-5
-10
10
15
20
25
r/nm
Figure 4.6. DLVO potential for various concentrations of NaCl (r = 12 nm is
the smallest distance possible for particles of 6 nm radius).
The maximum interaction potential, Umax, serves as the barrier to the two particles
combining. To determine the maximum potential, we graphed the sum of the van der Waals
component and the repulsion component versus r, the distance between the nanoparticles’
30
centers, as seen in Figure 4.6. We assume that the potential maximum corresponds to the
activation energy. We then used the Arrhenius equation (Equation 4.7) to approximate the
rate coefficient at the maximum potential.
k 2 = A freq e −U max / kT
(4.7)
The frequency factor, A freq , is calculated using the standard diffusion theory result
seen in Equation 4.8. It depends on the diffusion of the silver nanoparticles through an
aqueous solution. The rate constant for a diffusion-controlled reaction is: k2 = Afreq. For a
diffusion-controlled reaction between two identical particles, the diffusion rate constant is:
k d = 8πR * DN A , where R* is the encounter radius and D is the diffusion constant coefficient
of the nanoparticles in aqueous solution:23
D =
kT
6πηa
(4.8)
The hydrodynamic radius, a , is the effective radius in the solution taking into
account all the H2O molecules the particle carries in its hydration sphere. As the particles get
close, the hydrodynamic radii are equal to ½ R* resulting in Equation 4.9. This results in the
radii terms canceling out, making the reactive collision radius large; therefore, the particles
have a shorter distance to travel before colliding.
kd =
8πR ∗ N A kT
6πηa
Since R* = 2 a , k d =
A freq =
8(2a ) RT 8 RT
=
6ηa
3η
(4.9)
8RT
3η
With this approximation for the frequency factor, the rate constant is independent of
the identities of the reactants, and only depends on the temperature and the viscosity of the
solvent, η. For water at room temperature (298 K), the frequency factor is calculated to be
7.42 x 109 M-1s-1. Using the extinction coefficient for a 12 nm silver nanoparticle, 1 x 109 M1
cm-1, we can convert the frequency factor to 389 AU-1s-1. It is also important to note that as
the temperature of the reaction is varied, the viscosity of the water changes and therefore
alters the value of the frequency factor. Increasing the temperature increases the value of the
31
frequency factor and therefore increases the rate coefficient at the maximum potential, as
shown in Equation 4.7. Also, the viscosity of the water increases with decreasing
temperatures, therefore decreasing A freq . At this point all parameters needed to calculate k2
have been obtained except for the surface charge, Z.
The particle charge, Z, can be adjusted to obtain an agreement between the calculated
and observed values of k2 as a function of NaCl concentration. The experimental and
theoretical results are compared in Figure 4.7 for the three values of Z. For the onset of
aggregation of sodium chloride, a surface charge of 138 gave the best fit. However, there is a
decrease in the slope at concentrations above 32.5 mM. The drop of the slope may be
attributed to a charge increase on the surface of the nanoparticles and faster aggregation at
those concentrations (shown in Figure 4.7). The charge increases to 160. This is a result of
the nanoparticle volume enlarging, which results in a larger surface area and has a greater
capacity of obtaining a higher surface charge. At this point the surface of the nanoparticle is
not fully saturated by the sodium chloride; therefore, the nanoparticles aggregate faster. In
addition, passivation of the surface of the silver nanoparticle may influence the observed rate
constant and decrease the rate of aggregation.
Figure 4.7. Experimental and theoretical onset of aggregation for
NaCl.
32
4.5 DISCUSSION
Following the addition of sodium chloride to the silver nanoparticle solution, the
surface plasmon resonance peak increased in absorption and decreased in width by 15-30%.
This indicates that with the addition of sodium chloride, the surface of the silver nanoparticle
becomes modified. The chloride ions replaced the citrate and/or the oxide layer on the
surface of the silver nanoparticle. In Figure 4.1, we see the broad original plasmon band of
the silver nanoparticle solution due to the oxide layer of the silver nanoparticle. The decrease
in width is indicative of the chloride ions oxidizing and replacing the layer on the surface.
This modification of the surface of the nanoparticle allows for decomposition or aggregation
to initiate depending on the concentration of sodium chloride added.
Oxidative decomposition is the dominant process at sodium chloride concentrations
below 25 mM. At 1.5 mM a substantial decrease in the rate constant is observed due to
passivation. Passivation occurs when there is a build up of insoluble silver chloride that
covers the surface of the silver nanoparticle. This inhibits excess chloride ions to attach and
begin slow decomposition.
At higher concentrations of NaCl, the increasing ionic strength will elevate the degree
of charge screening and hence allow for an increase in aggregation kinetics. The rate of
reaction changes from 2/3rd order to 2nd order indicating silver nanoparticle aggregation. At
this point, the rate constants for the reaction increase tremendously. The functions of the
chloride ions, determined by the analysis using the DLVO Theory, were found to be that the
chloride ions not only increase the ionic strength of the solution, but also reduce the charge
on the surface of the nanoparticle. The aggregation behavior is consistent with the DLVO
Theory in that, at lower concentrations of sodium chloride, the increase of NaCl in solution
will increase the charge screening between the particles and allow an increase in aggregation.
At higher concentrations of NaCl, the charge on the surface of the silver nanoparticle is
completely screened and the potential barrier is eliminated, resulting in faster aggregation. At
this point, the kinetics have reached a maximum and have become independent of NaCl
concentration. This is the point were the critical coagulation concentration has been reached,
and the aggregation kinetics have become diffusion limited.
33
CHAPTER 5
THE ADDITION OF SODIUM FLUORIDE
Although fluoride is also a halide, it has a very different effect on silver nanoparticles.
Addition of low concentrations of sodium fluoride to silver nanoparticle solutions leaves the
surface plasmon resonance peak virtually unchanged. At higher concentrations, however,
NaF does induce aggregation, since it still affects the ionic strength, but the onset is higher
than that of the other halides.
5.1 SURFACE PLASMON RESONANCE PEAK
Unlike NaBr, NaI and NaCl, the addition of NaF to the nanoparticles does not alter
the shape of the SPR peak. As can be seen in Figure 5.1, there is no evidence of a decrease or
increase in absorption and/or change in the peak’s width and position. This indicates that
there is no fast absorption of the fluoride ion on the surface and that it does not replace the
citrate or oxide layer of the surface of the silver nanoparticle.
Figure 5.1. Changes in the surface plasmon resonance after NaF
addition (the final concentration of NaF is 10 mM).
Since there is no change in the surface plasmon resonance peak, we conclude that
there is little to no surface modification from the addition of the sodium fluoride. This is not
34
surprising given that AgF is much more soluble than the other halides, as shown in Table 5.1.
Therefore, we do not expect F- to complex with the silver but rather remain in the highly
favorable solvated state. Further evidence for this is given in the next section.
Table 5.1. Silver Halide Solubilities at 25°C
Silver Halide
Solubility (mol/100g of H2O)
AgF
1.397
AgCl
1.35 x 10-6
AgBr
7.18 x 10-8
AgI
1.11 x 10-9
5.2 OXIDATIVE DECOMPOSITION
Addition of NaF does not enhance the oxidative decomposition significantly. The
half-life of the SPR is almost as long as that when no halide is added. The rate of decay
remains extremely slow up to 80 mM NaF, at which point aggregation begins to take place.
The slow rate of decay is consistent with the fact that silver fluoride is very soluble and
associates more freely with H2O than the silver surface. Therefore, complexation of AgF
(Equation 5.1) does not cause the oxidative decomposition of the silver to become favorable.
Rather, the overall potential remains at the value of a reaction which does not produce a
complex (-0.3986 V):
1 / 2O2 + H 2 O + 2e − 
→ 2OH − ; E° = 0.401 V
2 Ag 
→ 2 Ag ( s ) + 2e − E° = -0.3986 V
;
(5.1)
5.3 SILVER NANOPARTICLE AGGREGATION
The onset of aggregation for the reaction of the silver nanoparticle solution with NaF
occurs at about 80 mM (see Figure 5.2). At this point, the rate becomes 2nd order. Using the
DLVO Theory discussed in Chapter 4, the same approach was taken to determine the surface
charge of the silver nanoparticle in the presence of NaF. From the concentrations of the NaF
35
Figure 5.2. Experimental and theoretical onset of aggregation for NaF.
solutions added to the silver nanoparticles, the ionic strength was calculated to solve the van
der Waals component using Equations 4.5 and 4.6. Knowing the size of the nanoparticle to
be 12 nm in diameter, the surface charge of the silver nanoparticle in the presence of NaF
was calculated to be 250. Comparisons of the calculated and experimental values of the rate
constant for aggregation are given in Figure 5.2.
5.4 DISCUSSION
From the analysis of the onset of aggregation for reactions with sodium fluoride using
the DLVO Theory, the surface charge of the silver nanoparticle was determined to be around
250. In Chapter 4, we estimated the surface charge of the silver nanoparticles reacting with
sodium chloride to be 138-160. Why is there a difference? To answer this we began by
comparing the surface plasmon resonance peaks for both. In the case of sodium chloride, it
was apparent that there was a significant change in the peak, indicating that the citrate and
oxide surface was being replaced by the chloride ions. In the case of sodium fluoride, there
was no evidence of any change indicating that the surface of the nanoparticle was not being
modified. We can then hypothesize that the initial surface charge of the silver nanoparticle is
250 and that the chloride ions reduced the charge on the surface to 138. Originally, the
surface of the silver nanoparticle is covered with sodium citrate, and as illustrated in Figure
5.3 the citrate has two negative charges, and when the chloride ion replaces the citrate, the
two negative charges are replaced by single negative charges. In order test our hypothesis, we
36
Figure 5.3. Replacement of citrate by Cl- on the surface of the silver
nanoparticles.
designed an experiment where we added a small amount of sodium chloride followed by the
addition of sodium fluoride. The experiment is discussed in Chapter 6.
37
CHAPTER 6
THE ADDITION OF SODIUM CHLORIDE
FOLLOWED BY SODIUM FLUORIDE
After analyzing the reactions of the silver nanoparticles with sodium chloride and
sodium fluoride, we hypothesized that the chloride ions modify the surface of the
nanoparticle, whereas the fluoride ions do not. This can be seen by the changes of the surface
plasmon resonance peak of the nanoparticles with the chloride ions and the lack of changes
with the fluoride ions at low concentrations. A small amount of chloride replaces the citrate
on the surface, and lowers the charge, but not enough to aggregate. The fluoride does not
change the surface, but cause aggregation. To show that the chloride ions function not only to
increase the ionic strength of the solution, but also to modify the surface of the nanoparticle,
we added a small amount of sodium chloride followed by various amounts of sodium
fluoride.
6.1 EXPERIMENTAL
To initiate the experiment, a 10 mM solution of sodium chloride was added to the
silver nanoparticles. A concentration of 10 mM was chosen because the concentration is low
enough to easily monitor the rate of decay and significantly low enough in concentration to
ensure that no particle aggregation can occur. After the sodium chloride is added, the
nanoparticles are allowed to react with the chloride ions for approximately two hours to make
certain that the surface of the nanoparticle is fully saturated and modified by the chloride
ions. The addition of 10 mM NaCl causes the surface plasmon resonance peak to
immediately sharpen and increase in absorbance (as discussed in Chapter 4).
After the solution had reacted for two hours, allowing the surface modification to
approach completion, different amounts of NaF were added. The reactions were monitored,
using the HP 8452A UV-Visible spectrophotometer, as a function of time. These experiments
were repeated several times.
38
In Chapter 5, the kinetics for the reactions of the silver nanoparticles with sodium
fluoride showed the onset of aggregation occurs at a concentration of 80 mM. However,
reactions between the silver nanoparticles with NaCl showed an earlier onset of aggregation
at a halide concentration of 25 mM. NaF and NaCl both increase the ionic strength of the
solutions, but NaCl causes aggregation much earlier than NaF. In order to understand why
there is such a drastic difference in the onset of silver nanoparticle aggregation, the
proceeding experiment was designed.
6.2 RESULTS AND DISCUSSION
From the kinetic rate law equations discussed in Chapter 3, the second order constants
for the reactions were calculated. The averaged rate constants as a function of salt
concentrations are shown in Figure 6.1. Also included in the figure are the results obtained
when NaCl alone was added to induce aggregation. There is a significant overlap of these
two sets of results. Thus, following the addition of NaCl with NaF resulted in rate constants
similar to those of NaCl. Adding the sodium chloride modified the surface enough to allow
the fluoride ions to passivate the surface of the silver nanoparticle and caused aggregation to
occur earlier than with NaF alone. This resulted in a shift of the onset of aggregation for NaF
from a concentration of 80 mM to 25 mM.
Figure 6.1. Rate constants for the reaction with 10 mM NaCl and NaF.
39
The change in the onset of aggregation indicated that the functions of the chloride
ions include: reduction of the charges on the surface and modification of the surface of the
silver nanoparticle. Because sodium fluoride does not change the surface of the silver
nanoparticle, the original surface charge of the silver nanoparticles is approximately 250.
When the DLVO Theory is used to analyze the NaCl data, a decrease of the surface charge
from 250 to 138 is observed, indicating that the NaCl reduces the surface charge of the
nanoparticle.
40
CHAPTER 7
TEMPERATURE STUDIES
It is important to study the stability of silver nanoparticles under varying
environmental conditions. A reduction in the stability of the silver nanoparticles can result in
aggregation. The aggregation of the silver nanoparticles can alter their characteristics and can
lead to different conclusions and observations in research. The most important environmental
factor that affects the stability, reactivity and characterization of silver nanoparticles is
temperature 24. When synthesizing silver nanoparticles, a change in the temperature results in
changes in the shapes of the silver nanoparticles. In this study, we did not alter the
synthesizing temperature, but we changed the temperature at which the decay reactions with
NaCl occurred.
7.1 VARYING TEMPERATURES OF THE SILVER
NANOPARTICLES AND HALIDE SOLUTIONS
The decay reactions of the silver nanoparticles were performed at temperatures
ranging from 4°C to 80°C using an orbital/mixing dry bath. Since the sodium chloride decay
reactions are very reproducible, NaCl was studied in order to evaluate the temperature
dependence of aggregation. We used 40 mM sodium chloride solutions because this
concentration is past the onset of aggregation. At this concentration, the reaction is fast
enough to be complete in a few hours, but slow enough to monitor multiple measurements
for each experiment.
Silver nanoparticles were synthesized using the method described in Chapter 2,
section 2. After allowing the silver nanoparticles to age for at least four days, the solutions
were brought to the target temperature for the initiation of the decay reactions. Along with
the silver nanoparticle solutions, the sodium chloride solutions were also heated or cooled to
the desired temperature. Once the chosen temperature was reached, the sodium chloride was
added to the nanoparticle solution, initiating the decay reactions. As seen in the case of the
room temperature decay reactions, the absorbance spectra were collected as a function of
time.
41
In Figure 7.1 the natural logarithm of the inverse half-life of each reaction, which is a
measure of the reaction rate, is plotted as a function of temperature. For temperatures below
24°C, the aggregation rate is consistently greater than that at room temperature. The color of
the silver nanoparticles changes from bright yellow to an iridescent yellow-green color
indicating possible surface and/or structural changes of the nanoparticles. Evidence for
aggregation is shown in the UV-Visible spectrum of the decay reaction at 8°C (Figure 7.2).
Not only does the peak near 400 nm decrease, but a broad absorbance to the red also
develops. As the temperature is increased above 37°C, the rate increases, as expected for
Arrhenius like reactions with a positive activation barrier. Wei et al., who looked at silver
nanoparticle stability, found that the aggregation process is promoted by an increase in
temperature. However, in Figure 7.1, we see that at the lowest temperatures, the rate
constant for aggregation is at its highest, indicating a more complex mechanism at the lower
temperatures. This will be further explored in the subsequent section.
Figure 7.1. Varying temperature effect on the half-lives of the silver
nanoparticles.
42
Figure 7.2. Aggregation at 40 mM NaCl at 8°C.
To determine if the observed rates were due to the reactions with the halides at the
specific temperature or due to changes on the surface of the nanoparticle caused by the
temperature, a solution of pure silver nanoparticles was heated to 65°C, and the UV-visible
absorption spectrum was monitored with time. The surface plasmon resonance peak
absorbance increased in time up until 223 minutes, at which point the absorbance began to
level off. Figure 7.3 shows the increase in the surface plasmon peak area and the leveling off
at an absorbance of 1.28. We interpret this as an increase in the rate of surface oxidation
caused by the heating process, such that complete oxidation of the surface occurred more
rapidly.
The next temperature studied was 40°C. Unlike what was observed at 65°C, the
absorbance of the surface plasmon peak did not increase as much after four days of heating
(Figure 7.4), indicating little modification of the surface of the silver nanoparticle. Since the
decay reactions were also carried out at temperatures below 24°C, a solution of silver
nanoparticles was cooled to 8°C, in order to observe if the surfaces were also changing
(Figure 7.4). No change in the SPR was observed. These studies showed that after a certain
temperature, the surface of the silver nanoparticles may become modified prior to halide
addition, thus influencing the rate of aggregation. Therefore, using the data for temperatures
above 45°C may affect our analysis and because of that, they have not been included in this
discussion.
43
Figure 7.3. Heated silver nanoparticle solution at 65°C. A) Increase in
the surface plasmon resonance peak after time, and B) the change in
the SPR peak area.
44
Figure 7.4. A) Heated silver nanoparticle solution at 40°C after 4
days. B) Cooled Silver Nanoparticles at 8°C after 115 minutes.
45
7.2 POSSIBLE MECHANISMS
To analyze the data, the Arrhenius equation was used. The Arrhenius equation gives
the dependence of the rate constant, k, of chemical reactions at temperature, T.
k = Ae − Ea / RT
(7.1)
The equation implies that for positive activation energy, as the temperature decreases,
the rate is expected to decrease. However, as discussed above, the opposite was observed. In
Figure 7.5, ln k vs. 1/T is plotted for temperatures below 45°C. The values of k were
calculated from the slopes of the plots of 1/A vs. time (See Chapter 4). As can be seen in the
Figure 7.5, there is a steady decrease in the rate as the temperature increases, implying a
negative activation energy. Although unusual, negative activations do occur when a reaction
is preceded by a fast equilibrium.
Figure 7.5. Ln k vs. 1/T, results in an activation of -6.561 x 104 J/mole.
A linear least squares analysis of the data, shown in Figure 7.5, results in a value of
the slope which is related to the activation energy. A value of -6.561 x 104 J was obtained.
A negative activation energy indicates the presence of a barrierless reaction, in which the
reaction proceeding relies on the capture of the molecules in a potential well 23. When the
temperature increases there is a reduced probability of colliding molecules capturing one
another. The cross section decreases with increasing temperature, and therefore the rate of
aggregation decreases.
46
For a reaction with a pre-equilibrium, there are three activation energies to take into
account. There are the forward and reverse steps of the pre-equilibrium and the final step. In
order to obtain an overall negative activation energy, the relative activation energies of the
steps are constrained. One possible arrangement is shown in the potential energy diagram
(Figure 7.6), where the sum of the forward step of the pre-equilibrium activation energy,
Ea(a), and the final step, Ea(b), is smaller than the reverse step of the pre-equilibrium
activation energy, Ea’(a). Therefore, the system exhibits negative temperature dependence,
where the rate decreases as temperature increases.23
Potential energy
Ea (a)
Ea ‘(a)
Ea (b)
A+A
∆H°
A2
A--A
Reaction coordinate
Figure 7.6. Potential diagram with a
negative activation energy.
In order to more clearly understand the mechanism and how the rate changes with
temperature, the bimolecular aggregation reaction must be further studied. We propose that
the bimolecular reaction forms an intermediate that consists of loosely bound nanoparticles,
A--A, which later form dimers, A2, which eventually aggregate further. The overall reaction
and the analysis of its rate law are given below:
k1
k2
A + A ⇔ A ⋅ ⋅ A →
A2
k−1
where: rate = k 2 K [ A]2 = k eff [ A] 2
47
K=
k1
= e − ∆ r G ° / RT = e − ∆H ° / RT + ∆S ° / R = e ∆S ° / R e − ∆H ° / RT
k −1
(7.2)
k 2 = Ae − Ea / RT
So, keff = Ae ∆S ° / R e − ( Ea+∆H °) / RT
Therefore, rate = Ae ∆S ° / R e − ( Ea+∆H °) / RT [ A]2
We can determine the temperature dependence of the system by looking at the Gibbs
free energy. Since this is a favorable reaction, the Gibbs free energy is negative. By using the
Gibbs free energy equation, K can be estimated in terms of enthalpy and entropy. Therefore,
keff can be evaluated in terms of the entropy, enthalpy and activation energy to observe any
temperature dependency. Since this system consists of a negative composite activation
energy, and a negative enthalpy, ∆H°, for a favorable reaction, the sum of the enthalpy and
activation energy remains negative as long as ∆H° is more negative than Ea. Therefore, the
system shows a negative temperature dependency.
This means that the pre-equilibrium is very fast, and that at lower temperatures the
formation rate of A--A is higher, producing more A2. As the temperature is increased, the
equilibrium shifts to the left, depleting the A--A complex and thus the rate of formation of A2
decreases.
7.3 HEATING SILVER NANOPARTICLES BEFORE
INITIATING DECAY REACTIONS
It has already been shown that a modification of the surface of silver nanoparticles
can lead to faster aggregation. Although, the bimolecular mechanism described in the
previous section is very plausible, the surface of the nanoparticle can not be ignored. In
Chapter 4, it was determined that the sodium chloride not only increased the ionic strength of
the solution, but also reduced the surface charge, resulting in an earlier onset of aggregation.
At temperatures below 37°C, it is also possible that the surface charge decreases, therefore
increasing the rate of aggregation.
We explored the possibility of a temperature-dependent surface charge by preheating
the silver nanoparticle solutions at 65°C and then measuring the decay reaction induced by
sodium fluoride. If the surface charge changes upon heating, a different aggregation onset
concentration would be expected. Since the SPR peak absorbance changed as the solutions
48
were heated at 65°C, it was important to determine whether the change in the peak altered the
previous results. This was further explored with the addition of sodium fluoride.
To begin the decay reaction, the silver nanoparticles were heated for 180-300 minutes
to ensure that the surface plasmon band had leveled off. The nanoparticle solutions were then
cooled to room temperature. The sodium fluoride mixture was added and the reaction was
monitored using the UV-visible spectrophotometer. The reaction half-lives were determined,
and the results are plotted in Figure 7.7, along with the results described in Chapter 5 (no
preheating of the nanoparticle solutions). The rate appears to be higher for the preheated
samples, indicating that the silver nanoparticle’s surface is slightly more oxidized. The
higher rate suggests that the particles have a lower surface charge. On the other hand, the
onset of aggregation does not appear to be significantly different, suggesting that the change
in the surface charge is not large.
Figure 7.7. Addition of NaF following the heating of the silver
nanoparticles.
The decay reactions of the silver nanoparticles exhibited a surprising temperaturedependent behavior. At low temperature, where one would expect a lower rate, the silver
nanoparticles exhibited anti-Arrhenius behavior, in which the rate of aggregation was highest
at the lower temperatures. As the temperature was increased to 37°C, the rate of aggregation
decreased to its lowest value. Above 37°C, the rate followed Arrhenius behavior, where the
rate increased with increasing temperatures. This indicates that the mechanism of the
49
aggregation of silver nanoparticles is complex and suggests that the first step in the reaction
is almost barrierless, where the molecules are captured in a potential well.23
It was also determined that the surface plasmon resonance peak did not change at
cooler temperatures. However, at 65°C the absorbance and the area of the SPR increased by
approximately 42%. At a temperature of 40°C, an intermediate change was observed, the
absorbance and the area increased by approximately 4%. By preheating the silver
nanoparticles prior to initiating the decay reactions, an increase of the rate of aggregation was
observed (as shown in Figure 7.7), but the onset of aggregation for NaF remained at 70 mM.
50
CHAPTER 8
CONCLUSIONS
With the increasing use of the silver nanoparticles in research and in industrial
applications, it is important to understand how silver nanoparticles react with the
environment. The goal of this research was to acquire a more in depth understanding of how
the silver nanoparticles react under various chemical and physical conditions. The study
began with a preliminary literary research to investigate the plausible kinetic processes of the
silver nanoparticles with the addition of sodium halides. The different sodium halides studied
were sodium bromide, iodide, fluoride and chloride.
The reaction between the silver nanoparticles and sodium bromide resulted in small
changes in the surface plasmon peak. The absorption peak exhibited a red shift, a slight
increase in the peak absorbance and a small decrease in width. At low concentrations of
sodium bromide, the reaction followed a 2/3rd order oxidative decomposition up until a 50
mM, where it changed to a 2nd order aggregation. In oxidative decomposition, the
nanoparticles lose silver atoms because of oxidation occurring on the surface. At high
concentrations, aggregation was observed. In this process, the slow step is the combination of
two particles, and the loss of silver atoms is due to the loss of silver monomers in the
solution.
The second halide studied was sodium iodide. Similar to the reaction with sodium
bromide, at low concentrations the reactions followed a 2/3rd order oxidative decomposition.
However, the order of the reaction remained 2/3rd even at high concentrations above 50 mM.
The addition of I- did not show any indication of forming a layer of passivation on the
surface of the silver nanoparticle. The surface plasmon peak exhibited a greater change than
that with Br-, a broadening of the peak and a substantial decrease in the absorption. The peak
also exhibited a significant blue shift. As the reaction reached equilibrium and/or completion,
a new peak at approximately 428 nm appeared. This indicated the presence of AgI. These
changes in the SPR peak indicate a significant change of the surface of the nanoparticles.
51
The addition of sodium chloride changed the kinetics and the surface plasmon
resonance peak, depending on the concentration added to the silver nanoparticles. The
surface plasmon band decreased in width and increased in absorbance. This indicated that the
chloride ions were adsorbing on the surface of the silver nanoparticle. The excess of chloride
ions also replaced the citrate and the oxide layer on the surface, resulting in the observed
change in the UV-Visible spectrum. At concentrations below 3 mM of sodium chloride, a
2/3rd order oxidative decomposition was dominant. This led to an increase of passivation on
the surface. A thick layer of AgCl accumulated on the surface of the silver nanoparticles, and
at concentrations between 3 mM and 25 mM no reaction was observed due to the layer on the
surface. At concentrations above 25 mM the rate order changes to 2, indicating the onset of
aggregation.
Unlike NaCl, NaF does not enhance oxidative decomposition of the silver
nanoparticles. The half-life remains consistently longer at all concentrations ranging from 0
to 60 mM. The fact that NaF is highly solvated in aqueous solution can explain why no
decomposition is observed. The complexation of AgF is not favorable and has an
electrochemical potential of -0.3986 V. The absence of oxidation and the lack of changes in
the surface plasmon peak indicate that there is no modification on the surface of the silver
nanoparticle. The onset of aggregation occurs at around 75 mM, at which point the rate
increases. Using the DLVO theory, we estimated the surface charge was higher for NaF than
NaCl, 138 and 250 respectively; i.e., the chloride ions reduced the charge on the surface of
the silver nanoparticles.
In order to test our hypothesis, a small amount of NaCl was added to the silver
nanoparticle solution in order to cover the surface of the nanoparticles. After two hours, NaF
was added to the solution. As stated in Chapter 6, the NaCl reduced the charge on the silver
nanoparticles and modified the surface. This allowed the fluoride to aggregate and shifted the
onset of aggregation from 75 mM to 25 mM, where the onset typically occurs for NaCl. This
indicated that the chloride ions changed the surface of the silver nanoparticle and allowed the
fluoride ion to initiate aggregation.
Decay reactions initiated at various temperatures exhibited interesting results.
Aggregation reactions due to 40 mM NaCl were carried out at temperatures ranging from 480°C and were monitored in time. At low temperatures the rate constants were the highest.
52
As the reactions approached room temperature, the rate constants were almost two orders of
magnitude slower. The rate constants slowly increased at temperatures above 37°C. The
mechanism for the aggregation process at low temperature was determined to be more
complex, showing Anti-Arrhenius behavior at low temperatures. The system exhibited a
negative temperature dependence, however, the aggregation process follows normal
Arrhenius behavior as the temperature is increased above 37°C.
53
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