Reactions of Aluminum with Halogen Containing Oxides
by
Cory Farley, B.S.M.E, M.S.B.M.E.
A Dissertation
In
MECHANICAL ENGINEERING
Submitted to the Graduate Faculty
of Texas Tech University in
Partial Fulfillment of
the Requirements for
the Degree of
DOCTOR OF PHILOSOPHY
Approved
Michelle Pantoya
Chair of Committee
Kevin Long
Gordon Christopher
Jordan Berg
Sukalyan Bhattacharya
Dominick Casadonte
Interim Dean of the Graduate School
May 2013
Copyright 2013, Cory Farley
ACKNOWLEDGEMENTS
“For I know the plans I have for you,” declares the Lord, “plans to prosper you and not to
harm you, plans to give you hope and a future.” – Jeremiah 29:11. Rarely do we take a
direct path to the place God wants us to be, but the destination is worth it.
Dr Michelle Pantoya, thank you for the patience you have shown to me over the years as
I stumbled my way through. From the time I showed up at your office unannounced and
unexpected and you recruited me to the Combustion Lab, you have been a rock on which
I could rely. When I was confused on my research, lost in my classes, or panicking on my
thesis you were there to guide me back on my path. I only hope that I can portray your
professionalism and success as I move forward in my career.
To my family, I left for school a long time ago as a scared little kid. However, I have
always had your support to do anything, and this was no different. Dad, thank you for
instilling a work ethic. Mimi and Papa, thank you for the financial help. I may have eaten
ramen noodles for a stretch, but I have no student loans, and there is tremendous peace of
mind knowing I will be moving forward debt free.
To my beloved wife Stephanie and my son Eli, thank you for providing sanctuary. My
home has always been a place of serenity where I am free to love and be loved.
Stephanie, you are my soulmate with whom there are no words to describe how powerful
my feelings are for you and with Eli, you are my proof that God’s plan for me does
indeed give me “hope and a future”
Also, I would like to acknowledge Santanu Chaudhuri and Martin Losada at ISP/Applied
Sciences Laboratory, Washington State University, WA for their efforts running the
quantum mechanical modeling seen in Chapter III.
ii
TABLE OF CONTENTS
ACKNOWLEDGEMENTS ....................................................................................................... II
LIST OF TABLES ................................................................................................................. IV
LIST OF FIGURES ................................................................................................................ V
ABSTRACT .......................................................................................................................... VI
CHAPTER I: BACKGROUND ................................................................................................ 1
Ignition ............................................................................................................................ 1
Propagation ..................................................................................................................... 5
Flame Speed Measurement ........................................................................................... 12
Conclusion .................................................................................................................... 14
CHAPTER II: REACTION KINETICS OF NANOMETRIC ALUMINUM AND IODINE
PENTOXIDE ....................................................................................................................... 15
Introduction ................................................................................................................... 15
Experimental Method.................................................................................................... 16
Results and Discussion ................................................................................................. 17
CHAPTER III: REACTION KINETICS OF NANO-ENERGETICS WITH HALOGEN
CONTAINING OXIDES ....................................................................................................... 23
Introduction ................................................................................................................... 23
Methods......................................................................................................................... 24
Near Equilibrium Experiments ............................................................................................... 25
Non-Equilibrium Experiments ................................................................................................ 26
Ab Initio Quantum Chemical Calculations ............................................................................. 29
Solid State Calculations ........................................................................................................ 30
Gas-phase Calculations. ....................................................................................................... 31
Results ........................................................................................................................... 33
Discussion ..................................................................................................................... 39
CHAPTER IV: THE CONTRIBUTION OF ATMOSPHERIC OXYGEN TO ENERGY
PROPAGATION IN REACTIVE MATERIALS ....................................................................... 43
Introduction ................................................................................................................... 43
Methods......................................................................................................................... 46
Results ........................................................................................................................... 48
Discussion ..................................................................................................................... 50
Iodine pentoxide .................................................................................................................... 51
Copper oxide ......................................................................................................................... 52
Calcium Iodate....................................................................................................................... 52
Iron oxide ............................................................................................................................... 53
CONCLUSIONS ................................................................................................................... 55
BIBLIOGRAPHY ................................................................................................................. 57
iii
LIST OF TABLES
1. Powder characteristics .................................................................................................. 16
2. Reactant average particle diameter, supplier and anticipated products
from Al reaction with respective oxidizer .................................................................... 25
3. Summary of experimental results. ................................................................................ 36
4. Decomposition of iodine oxide clusters in gas phase for B3LYP, PBE0
and ACM functionals at aug-cc-pVTZ (aVTZ) and aug-cc-pVQZ
(aVQZ) levels of theory. The values in the bracket are calculated using
spin-orbit (SO) corrections. .......................................................................................... 37
5. The reaction enthalpy and free energy calculated for Al reactions with
iodine oxide fragments calculated at aug-cc-pVTZ (aVTZ) and aug-ccpVQZ (aVQZ) .............................................................................................................. 38
6. Gas production from REAL Code analyses ................................................................. 48
7. Summary of flame speed results in units of m/s. ......................................................... 49
iv
LIST OF FIGURES
1. Energy changes as an exothermic reaction proceeds .................................................... 2
2. Concept sketch of temperature profile across reaction zone of the
heterogeneous composite. ............................................................................................. 6
3. Heat flow and mass loss curves of I2O5; experiments performed in an
Ar environment at a heating rate of 10°C/min. ........................................................... 17
4. Heat flow and mass loss curves for Al2O3 and I2O5, experiments
performed in an Ar environment at a heating rate of 10°C/min. ................................ 18
5. Heat flow and mass loss curves for micron scale Al and I2O5,
experiments performed in an Ar environment at a heating rate of
10°C/min. Note the lack of any exothermic reaction .................................................. 20
6. Nanometer scale Al powder and I2O5 with an equivalence ratio of 0.4,
experiments performed in an Ar environment at a heating rate of
10°C/min. .................................................................................................................... 21
7. Schematic of bomb calorimetry experiment. .............................................................. 27
8. Schematic for open tray flame speed measurements. ................................................. 27
9. Heat flow measurements for Al+I2O5 scanned at 2, 5 and 10 K/min
heating rates. ............................................................................................................... 33
10. Peak temperatures for Al+I2O5 plotted using the Type B-1.95 Peak
Method. The resulting Ea is 175.6 kJ/mol .................................................................. 33
11. Graph showing Al+I2O5 Parr bomb calorimetry result displaying dT ...................... 34
12. Open tray still frame images A) Al+I2O5 taken at 33µs intervals, B)
Al+Ca(IO3)2 at 17ms intervals, C) Al+AgIO3 at 40 µs intervals, D)
Al+Fe2O3 at 54ms intervals. ....................................................................................... 35
13. Position of flame front as a function of time for one test for Al+I2O5
with a flame speed of 898 m/s. ................................................................................... 35
14. Top view schematic for open tray flame speed experiments. ..................................... 48
15. A) Al+Ca(IO3)2 at 6.9ms intervals. B) Al+Fe2O3 at 6.9ms intervals. C)
Al+I2O5 at 44µs intervals. D) Al+CuO at 121µs intervals. ........................................ 49
16. Flowchart of fast convective reactions ....................................................................... 51
17. Percent increase in flame speed as a function of sensitivity factor ............................. 54
v
ABSTRACT
Due to increasing threats of biological attacks, new methods for the neutralization
of spore forming bacteria are currently being examined. Thermites may be an effective
method to produce high temperature reactions, and some compositions such as aluminum
(Al) and iodine pentoxide (I2O5) also have biocidal properties. This study examines the
thermal degradation behavior of I2O5 mixed with micron and nanometer scale aluminum
(Al) particles. Differential scanning calorimetry (DSC) and thermo-gravimetric (TG)
analyses were performed in an argon environment on both particle scales revealing a nonreaction for micron Al and a complex multistep reaction for the nanometer scale Al.
Results show that upon I2O5 decomposition, iodine ion adsorption into the alumina shell
passivating Al particles is the rate controlling step of the Al-I2O5 reaction. This preignition reaction is unique to nano-Al mixtures and attributed to the significantly higher
specific surface area of the nanometric Al particles which provide increased sites for Isorption. A similar pre-ignition reaction had previously been observed with fluoride ions
and the alumina shell passivating Al particles.
Composite energetic materials comprised of nanoparticle fuel and oxidizer can
exhibit high flame propagation speeds on the order of 1000m/s when burning in an
unconfined environment. In particular, halogen based formulations such as aluminum and
iodine pentoxide have received significant attention due to both high flame speeds and
biocidal properties. Studies have attributed high flame speeds to convective influences
within a reaction driving the heat forward in a pressure wave accelerating ignition of
unburned powders. This study examines factors contributing to convective flows such as
gas and heat generation and their relationship to the measured flame speed as well as
fundamental chemical kinetics influencing the observed flame speeds. The goal is to
understand parameters directly related to high flame speeds in halogen containing
composites. Results show a direct correlation between apparent activation energy and
flame speed indicating that flame speed is directly influenced by chemical kinetics. For
vi
this reason, the intermediate chemistry associated with Al and iodine species was
examined to identify chemical influences accelerating flame speeds. Ab initio quantum
chemical calculations of gas-phase reactions resolve key exothermic intermediate
reactions contributing toward the kinetics of the fastest burning mixtures. Condensed
phase density functional theory calculations of Al2O3/I2O5 interfaces resolved
adsorption/desorption properties.
This study examines the effect of atmospheric oxygen concentration (4 or
93% oxygen) on energy propagation of nanometric aluminum with copper oxide
(Al+CuO), iron oxide (Al+Fe2O3), calcium iodate (Al+Ca(IO3)2), and iodine pentoxide
(I2O5). In all cases energy propagation was examined in terms of flame speed and higher
in the high oxygen environments. However, the convectively driven mixtures showed a
smaller percent increase in flame speeds. This behavior is attributed to the increased
availability of oxygen as a result of solid oxidizer thermal decomposition at lower
temperatures. The slower Al+Fe2O3 reaction showed greater increases in flame speed
attributed to early stage reactions involving atmospheric oxygen that promote oxide
decomposition and faster flame speeds. A predictor based on solid oxidizer thermal
decomposition and flame speed was developed to evaluate the sensitivity of a reaction to
atmospheric oxygen concentration.
vii
CHAPTER I: BACKGROUND
During the scientific revolution of the 17th century, Johann Becher in his Physica
Subterranea (1667), claimed that the minerals of the earth were made of “terra fluida”
(fluid earth), terra phingus (fatty earth), and terra lapidea (stony earth) (Bowler and
Morus 2005). According to Becher, the terra phingus content made a material flammable.
In the early 17th century, Georg Ernst Stahl continued Becher’s work and renamed terra
phingus “phlogiston.” Combustion occurred when a material’s phlogiston was released,
and a substance such as ash, could not burn because it no longer contained phlogiston.
The theory of phlogiston would hold until the 1770’s with the discovery of oxygen. The
discovery of oxygen can be attributed to either Carl Scheele and his “fire air” (1772) ,
Joseph Priestly and his “dephlogisticated air” (1775) or Antoine-Laurent Lavoisier who
coined the term “oxygène” (1775). Earlier in 1772, Lavoisier had claimed that
combustion was a combination of burning substance and air. Later upon reading about
Priestly’s dephlogisticated air, Lavoisier refined his theory and argued that it was the
dephlogisticated air, which he labeled “oxygène” that contributed to combustion, and the
modern concept of combustion was born (Bowler and Morus 2005).
According to Merriam Webster Dictionary, combustion is defined as “a usually
rapid chemical process (as oxidation) that produces heat and usually
light.”(“Combustion”) For combustion to take place, the reactants must come together
with enough energy to initiate just as any chemical reaction. In terms of combustion, the
reactants are usually a fuel (methane, carbon, aluminum) and an oxidizer (oxygen gas,
metal oxide, halogen gas). Once initiated, if the energy transferred to the surrounding
reactants reaches a certain level, then initiation and transfer take place again and again,
and the combustion propagates. The continued propagation of the chemical reaction
results in a flame that was studied for burn rate and heat production in this paper.
Ignition
The first step for combustion is the initiation of the chemical reaction achieved
when a certain energy threshold for the reactants has been met. Svante Arrhenius studied
this energy threshold and coined the term “activation energy,” which is visualized in
Figure 1 where 𝐸𝑎 denotes activation energy, ∆𝐻𝑐 heat of combustion.
1
Figure 1. Energy changes as an exothermic reaction proceeds
Arrhenius went on to quantify the relationship between 𝐸𝑎 and reaction rate,
𝑘(𝑇)as seen below in Eq. (1),
𝑘(𝑇) = 𝐴 exp (−
𝐸𝑎
)
𝑅𝑇
(1)
where 𝐴 is a pre-exponential factor, 𝑅 is the gas constant, and 𝑇 is absolute
temperature.
As seen in Figure 1, when the activation threshold is met, the reaction progresses.
Combustion reactions are exothermic by definition, therefore the ensuing energy release
is greater than the activation energy and heat is released into the surroundings. If the
energy transferred to the neighboring reactants is greater than the activation energy, then
the chemical reaction initiates, and the combustion of the particles propagates along the
powder (Turns 2012).
The activation and initiation of a combustion reaction is called ignition.
Glassman describes two types of ignition: thermal and spark ignition (Glassman and
Yetter 2008). Thermal ignition can be active (i.e. a match) or passive resulting in a
spontaneous combustion with no active trigger. For example, a large pile of leaves can
spontaneously ignite as heat from the decaying process builds up. Simply put, as long as
iv
the rate of heat generation is greater than the rate of heat dissipation, the ensuing energy
buildup could ignite a combustion reaction. In 1935, Semenov modeled this behavior by
combining an energy balance with the Arrhenius form of the chemical rate equation to
develop a time constant comparison of the generation and dissipation energy terms as
seen in Eq (2).
𝑄
𝑒 (𝑐 )
𝜏𝑟
𝑣
( )=
2
𝑅𝑇
𝜏𝑙
𝑜
𝐸𝑎
(2)
where 𝜏𝑟 and 𝜏𝑙 are the time constants for the generation and dissipation
respectively, 𝑄 is the thermal energy release of the chemical reaction, 𝑅 and 𝑐𝑣 are the
ideal gas constant and specific heats, 𝐸𝑎 is activation energy, and 𝑒 is the number 𝑒.
For a more controlled ignition, a trigger is often used. For example,
gasoline is typically ignited with a spark in an automobile internal combustion engine.
Spark ignition is popular because it is reliable and safe due to the fact that a pilot light is
not required to initiate combustion. Most discussion on spark ignition centers on the term
“minimum ignition energy”. Minimum ignition energy describes the smallest energy
impulse that can ignite a given mixture under given circumstances. In order to determine
the minimum ignition energy required by a spark to initiate combustion, the spark itself is
often modeled as a sphere of hot gas. Using the quench criteria that the heat generated by
the reaction (or in this case spark) must be approximately equivalent to the rate of heat
loss of a volume via conduction, a minimum energy input for a spark, and hence, a gap
and voltage can be calculated. To meet the quenching criterion, Eq.(3) is introduced.
𝑄̇ ‴𝑉 = 𝑄̇𝑐𝑜𝑛𝑑
(3)
In Eq. (3), 𝑄̇ ‴ is the rate of heat generation, 𝑉 is the spherical volume and 𝑄̇𝑐𝑜𝑛𝑑
refers to the heat loss rate via conduction. The heat generated can be related to the
consumption of fuel (𝑚̇𝐹‴ ) by 𝑄̇ ‴ = −𝑚̇𝐹‴ Δ𝐻 where 𝐻𝑐 is the heat of combustion. The
heat lost to conduction can be modeled by Fourier’s law for an axisymmetric sphere as
seen in Eq (4)
iv
4 3
𝑑𝑇
2
−𝑚̇𝐹‴ Δ𝐻𝑐 ∙ 𝜋𝑅𝑐𝑟𝑖𝑡
= −𝜆4𝜋𝑅𝑐𝑟𝑖𝑡
( )
3
𝑑𝑟 𝑅𝑐𝑟𝑖𝑡
(4)
for a material with thermal conductivity 𝜆 and minimal spherical radius 𝑅𝑐𝑟𝑖𝑡
corresponding to the sphere representing minimal ignition energy. Substituting
−
𝑇𝑏 −𝑇∞
𝑅𝑐𝑟𝑖𝑡
=
𝑑𝑇
𝑑𝑟 𝑅 𝑐𝑟𝑖𝑡
and (𝜈 + 1)𝑐𝑝 (𝑇𝑏 − 𝑇∞ ) = Δ𝐻𝑐 for a given boundary temperature (𝑇𝑏 )
resulting in ignition in Eq. (4) allows for the determination of 𝑅𝑐𝑟𝑖𝑡 as seen in Eq (5).
𝑅𝑐𝑟𝑖𝑡 =
√6𝛼 √6
=
𝛿
𝑆𝐿
2
(5)
where 𝛼 is thermal diffusivity, and 𝑆𝐿 is the latent flame speed and 𝛿 is the
laminar flame thickness. Equation (5) indicates that the critical radius is roughly
equivalent to the laminar flame thickness. From the minimum radius (𝑅𝑐𝑟𝑖𝑡 ), we can
calculate a minimum mass (𝑚𝑐𝑟𝑖𝑡 ) and thus, a minimum energy (𝐸𝑖𝑔𝑛 ).
𝐸𝑖𝑔𝑛 = 𝑚𝑐𝑟𝑖𝑡 𝑐𝑝 (𝑇𝑏 − 𝑇∞ )
(6)
Recall that for a spark, the energy is
1
𝐸 = 𝑐𝑓 (𝑉22 − 𝑉12 )
2
(7)
where 𝑐𝑓 is the capacitance in farads, and 𝑉2 , 𝑉1 are the voltages after and before
the spark.
The ignition of binderless solid particles results from a complex series of
events adapted from a description by Kuo: (Kuo 1986)

Thermal energy transfer to particles via conduction, convection, radiation

Inert heating within the solid particles

Oxidizer phase changes and decomposition

Subsurface chemical reactions

Diffusion of fuel/oxidizer species

Heterogeneous reactions between gas/condensed phases
iv

Gas-Phase reactions

Abrupt temperature increase

Light emission

Development of a flame
According to a review in 1966 by Price, ignition theories are classified into three
major groups, 1) gas-phase, 2) heterogeneous, and 3) solid-phase (Goyne et al. 1966). In
the case of aluminum combustion, most reactions are believed to fall in the latter two
groups with nanometer scale aluminum particles having been shown to react in the solid
phase as described by (change in thermal energy) = (conduction) + (heat generation of
chemical reactions) + (energy accumulation due to in depth radiation absorption).
Reaction rates for these reactions are dependent on the chemical reaction rate
constant, 𝑘(𝑇), and how strongly the oxidizer adsorbs to the surface of the fuel (Turns
2012). If the oxidizer is weakly adsorbed, then the reaction rate, R, is proportional to the
concentration of oxidizer [O].
𝑅 ∝ 𝑘(𝑇)[𝑂]
(8)
Secondly, if the oxidizer is strongly adsorbed, then 𝑅 will depend mainly on 𝑘(𝑇)
𝑅 ∝ 𝑘(𝑇)
(9)
The remaining scenario involves a case where the combustion products
𝑃 compete with the oxidizer for adsorption sites on the surface of the fuel.
𝑅 ∝ 𝑘(𝑇)
[𝑂]
[𝑃]
(10
)
Propagation
After the flame has been ignited, the chemical reaction generates enough
energy to drive the surrounding reactants to their activation energy and ignition. The
“ignition and energy spread” cycle repeats with the surrounding reactants. The cyclic
repetition of 1) activation, 2) ignition and 3) energy transfer describes the concept of
flame propagation, which is displayed graphically in Figure 2.
iv
Figure 2. Concept sketch of temperature profile across reaction zone of the heterogeneous
composite.
Classically, the study of flame propagation has been focused on three theories as
summarized by Kuo.(Kuo 1986) He divided the theories into three major catergories:
thermal, comprehensive, and diffusive. The theories cited from 1883 to 1954 are listed
below with their corresponding citations.

Thermal Theory

Mallard and Le Chatelier’s Development

The Damkohler theory

The theory of Bartholome

The theory of Emmons, Harr and Strong

The theory of Bechert

Comprehensive Theory

The theory of Lewis and von Elbe

The theory of Zel’dovich, Frank-Kamenetsky, and Semenov
iv

The theory of Boys and Corner

The theory of Hirschfelder and Curtiss

The theory of von Karman and Penner

Diffusion Theory

The theory of Tanford and Pease

The theory of Van Tiggelen

The theory of Manson

The theory of Gaylord and Wolfhard
The purpose of laminar flame theories such as those listed above is to determine
laminar flame speed (𝑆𝐿 ). In 1883, Mallard and Le Chatelier proposed the first thermal
theories for flame propagation. Based on the temperature profile in Figure 2, they
proposed that the preheating was driven by heat conduction from the combustion zone of
the neighboring material. The flame speed is determined via energy balance and is found
to be proportional to the exponential term for the reaction rate equation (Eq. (1)) as seen
in Eq. (11) below.
𝑆𝐿 ∝ √𝑒
𝐸
− 𝑎
𝑅𝑢 𝑇
(11)
𝐸𝑎 is activation energy, 𝑅𝑢 is the molar ideal gas constant and 𝑇 is the temperature
at which the reaction takes place.
Later, Zeldovich, Frank-Kamenetskii, and Semenov would introduce the
concept of mass diffusion to the thermal theory of Mallard and Le Chatelier. In the new
comprehensive theory, mass diffusivity is assuming to be equal to thermal diffusivity
(𝒟 = 𝛼). This assumption leads to a combination of continuity and energy equations that
when solved leads to the flame speed equation seen in Eq. (12).
2
𝜆𝑐𝑝f
𝐸𝑎
−
𝑅𝑇
𝑆𝐿 ∝ √
3 𝑎𝑜 𝐴𝑜 𝑒
𝜌𝑜 (𝑐̅𝑝 )
iv
(12)
where 𝑐𝑝f is the specific heat at the flame temperature, and 𝑐̅𝑝 is the average
specific heat between the initial and flame temperatures. 𝐴𝑜 , 𝐸𝑎 are the pre-exponential
term and activation energy from the Ahhrenius equation. The term 𝑎𝑜 refers to the initial
reactant density while 𝜌𝑜 refers to the initial overall density.
In our modern age of computing, fewer assumptions need to be made
since numerical solutions can be resolved with much less effort. Modern software such as
CHEMKIN uses a scheme by Kee et al solves for mass burn rate as an eigenvalue
problem and determines the flame speed from the mass consumption combined with the
unburned gas density.
To understand the mechanics of solid combustion, a brief discussion on
the combustion of carbon can help shed some light into the reaction kinetics. Unlike
aluminum, carbon is not encased in any sort of protective shell; therefore, when burned in
air, the surface of the carbon particle comes into direct contact with oxygen. Due to this
intimate contact, the burning of a solid carbon particle can be analogously modeled to a
burning fuel droplet. A straightforward discussion of the early part of this derivation is
not readily available, therefore I included it here.
First, examine a first law balance for the surface of the particle in a onedimensional Cartesian system. Eq. (13)
∑ 𝑄̇𝐶𝑉 − ∑ 𝑊̇𝐶𝑉 + ∑ 𝑚̇(ℎ + 𝑘𝑒 + 𝑝𝑒)
𝑖𝑛
𝑑𝐸𝐶𝑉
− ∑ 𝑚̇(ℎ + 𝑘𝑒 + 𝑝𝑒) =
𝑑𝑡
(13)
𝑜𝑢𝑡
where 𝑄̇𝐶𝑉 , 𝑊̇𝐶𝑉 are the heat and work boundary transfer rates for the control volume,
(ℎ + 𝑘𝑒 + 𝑝𝑒) are the enthalpy, kinetic and potential energy transfers associated with a
given mass flow rate 𝑚̇, and 𝐸𝐶𝑉 is the total energy of the control volume.
Assuming:

Burning is steady state

Potential and kinetic energies are negligible

𝑊̇𝐶𝑉 is negligible
iv

Bulk mass (𝑚̇) is a steady, single flow
Equation (13) simplifies to
[𝑄̇ ″𝐴]𝑥 − [𝑄̇ ″𝐴]𝑥+∆𝑥 + [𝐴𝑚̇″ℎ]𝑥 − [𝐴𝑚̇″ℎ]𝑥+∆𝑥 = 0
(14
)
where (″) denotes a flux over a given area 𝐴.
Dividing Eq. (14) by (−𝐴∆𝑥) yields
𝑄̇ ″𝑥+∆𝑥 − 𝑄̇ ″𝑥 𝑚̇″ℎ𝑥+∆𝑥 − 𝑚̇″ℎ𝑥
+
=0
∆𝑥
∆𝑥
(15
)
Then, take the limit as ∆𝑥 → 0,
𝑑𝑄̇ ″
𝑑ℎ
+ 𝑚̇″
=0
𝑑𝑥
𝑑𝑥
(16
)
From here, we need to take a detailed look at the heat transfer, 𝑄̇ ″. Until now, we
have been treating the mass flow, 𝑚̇″, through the control volume as if it was a single
component, steady flow. However, this is a reacting mixture and each species in the flow
is changing as combustion takes place inside the control volume. In order to account for
this, it is common practice to include the deviations of each species diffusion rates into
the heat transfer term. Therefore, the heat transfer term includes both Fourier conduction
and the enthalpies of species deviating from the bulk flow when diffusing through the
carbon particle as seen below in Eq. 17,
𝑄̇ ″ = −𝜆
𝑑𝑇
+ ∑ 𝑚̇″𝑖,𝑑𝑖𝑓𝑓 ℎ𝑖
𝑑𝑥
(17
)
where 𝑚̇″𝑖,𝑑𝑖𝑓𝑓 is the mass diffusion deviating from bulk values with corresponding
enthalpy, ℎ𝑖 and 𝜆 is the thermal conductivity.
Recall from Fick’s law of mass diffusion,
𝑚̇″𝑖,𝑑𝑖𝑓𝑓 = −𝜌𝒟
Substituting Eq. (18) into Eq. (17) yields
iv
𝑑𝑌𝑖
𝑑𝑥
(18
)
𝑄̇ ″ = −𝜆
𝑑𝑇
𝑑𝑌𝑖
− 𝜌𝒟 ∑
ℎ ,
𝑑𝑥
𝑑𝑥 𝑖
(19
)
assuming all the species have a similar diffusivity, 𝒟.
From here, following Turns (Turns 2012), it can be shown that the Shvab-Zeldovich
energy equation for spherical geometry is
𝜕 ∫ 𝑐𝑝 𝑑𝑇
1 𝑑 2
𝑜
[𝑟
(𝜌𝑣
∫
𝑐
𝑑𝑇
−
𝜌𝒟
)] = − ∑ ℎ𝑓,𝑖
𝑚̇𝑖‴ .
𝑟
𝑝
𝑟 2 𝑑𝑟
𝜕𝑟
(20
)
In order to get to this equation, the following assumptions have been made:

Steady State

No kinetic or potential energy effects

No control volume work (i.e. shaft or piston)

No radiation

𝐿𝑒 = 𝒟 = 1
𝛼
Assuming a quasi-steady burning, 𝑚̇𝑟 is a constant with respect to 𝑟 therefore
𝑑(𝜌𝜈𝑟 𝑟 2 )
=0
𝑑𝑟
(21)
Recall that 𝜈𝑟 is bulk flow velocity. Thus, with constant properties, Eq. (20) can be
rewritten as
𝑚̇𝑐𝑝𝑔 𝑑𝑇
𝑑 2 𝑑𝑇
(𝑟
)=
𝑑𝑟
𝑑𝑟
4𝜋𝜆 𝑑𝑟
(22)
𝑜
with the assumption of no reactions (evaporation only, ∑ ℎ𝑓,𝑖
𝑚̇𝑖‴ = 0) at the surface of
the droplet. From here, it is straightforward to follow Turns (Turns 2012) to arrive at the
diffusion droplet lifetime equation
𝑡𝑑 = 𝐷𝑜2 /𝐾
(23)
where 𝑡𝑑 is total burn time, 𝐷𝑜 is initial diameter and 𝐾 is an evaporation constant
calculated from material properties and temperature differences.
iv
Now, burn time can be examined for a kinetically driven reaction. The flux of fuel
at the droplet surface can be described as
𝑚̇𝑓″ =
𝑚̇𝑓
𝑑𝑟𝑠
= 𝜌𝑠
2
4𝜋𝑟
𝑑𝑡
(24)
where the subscript 𝑠 denotes the particle surface.
While the flux of oxidizer can be described by
𝑚̇𝑓″
= 𝑘𝑠 [𝑂]𝑠
𝑖
𝑚̇𝑜″ =
(25)
where 𝑖 is the mass stoichiometric index, and 𝑘𝑠 is the reaction rate constant and [𝑂]𝑠 is
the concentration of oxidizer on the surface. However, [𝑂]𝑠 is not readily known.
However, if the reaction is kinetically limited, it can be assumed that the concentration on
the surface is equal to the free stream concentration for oxidizer, or
[𝑂]𝑠 = [𝑂]∞
(26)
Combining Eqs. (24), (25), and (26) yields
𝑚̇𝑓″ =
𝑚̇𝑓
𝑑𝑟𝑠
= 𝜌𝑠
= 𝑖𝑘𝑠 [𝑂]𝑠 = 𝑖𝑘𝑠 [𝑂]∞
2
4𝜋𝑟
𝑑𝑡
(27)
which simplifies to
𝑑𝑟𝑠 𝑖𝑘𝑠 [𝑂]∞
=
𝑑𝑡
𝜌𝑠
(28)
Separate the variables then integrate and solve for time.
𝑡𝑘 =
𝜌𝑠 𝐷𝑜
2𝑖𝜌[𝑂]∞ 𝑘𝑠
(29)
Therefore, in kinetic limited reactions, the burn time is proportional to the droplet
diameter or
𝑡𝑘 ∝ 𝐷𝑜1
iv
(30)
Flame Speed Measurement
Flame speed is typically defined for gas phase reactions as “the velocity at
which the unburned gases move through the combustion wave in a direction normal to
the wave surface.” (Glassman and Yetter 2008) For example, in a flame tube, the flow
rate that holds the flame stationary can be directly related to the flame speed. However, in
multiphase media with varying particle geometries the concept of flame speed becomes
less clear. Flame speeds for multiphase combustion processing involving solid phase
reactants, flame speed will be defined here as “the velocity at which the flame front
consumes and propagates through the unburned particles.”
Burn rate experiments for solid particle combustion can either consist of
measuring the burn time for a single particle or burn rate for powders. Early experiments
for calculating the burn time of solid particles involved dropping them into a flat flame,
and in 1963, Friedman and Macek conducted this experiment with aluminum particles
(Friedman and Maček 1963). The light traces from the burning particles were recorded
and used to estimate burn times that correlated to 𝐷1.5 power for aluminum particles from
10 to 74 µm. Recall, that for purely diffusion driven reactions, the burn time should be ∝
𝐷2 . For purely kinetic reactions, the correlation would be ∝ 𝐷1 . So, a reaction of ∝
𝐷1<𝑥<2 would have a mixture of diffusion and kinetic elements in the combustion. More
recently, burn time experiments for single particles has been conducted in shock tubes.
Beckstead, in 2005, collected data from a variety of experiments for particles ranging
from 20 to 800µm and found a correlation of 𝐷1.8 (Beckstead 2005). Lynch et al
correlated 3-11µ1m particles at 𝐷0.3 (Lynch, Krier, and Glumac 2009). Finally in 2011,
Gill et al used a CO2 laser to ignite 10-14µm particles and correlated them to
𝐷0.8−0.9 (Gill, Badiola, and Dreizin 2010). The reaction kinetics involved causing the
burn times to fall out of the range from kinetic to diffusion controlled are not fully
understood at this time.
The powder flame speed measurements typically fall into three categories:
aerosol, open tray and closed tube. Aerosols containing solid particles are extensively
studied in an effort to understand dust explosions. Solid particle aerosols behave similarly
to the premixed flame previously discussed (Proust and Veyssiere 1988). Bunsen burner
style experiments with a constant flow are difficult to execute due to the difficulty in
iv
mixing the solid particles into the flow. Instead, the powder is typically dispersed
throughout a volume and then ignited. For volumes such as an open duct, the flame
position is measured with optical high speed imaging. For spherical closed volumes, the
flame speed can be related to pressure traces from pressure transducers within the vessel
(Pu et al. 2007) or derived from optical techniques (Bradley, Chen, and Swithenbank
1989).
For combustion of masses of non-aerosolized powders, flame speed
measurements are typically divided into open and closed measurements. For closed tube
measurements, powder is loaded into a tube in such a way as to maintain a constant
overall bulk density and ignited by a spark, hot wire or other stimuli (M. Weismiller et al.
2009; Bockmon et al. 2005). Optical readings can be taken by either high speed imaging
or photodiodes. High speed cameras give more information to help determine the nature
of the combustion and make accurate flame speed estimates. However, in many cases,
flame speeds are so high (~2000𝑚/𝑠) that a camera with the resolution and frame rate
required to get reliable information may prove cost prohibitive (Martirosyan, Wang, and
Luss 2009). Photodiodes give similar information at a much lower cost. A photodiode
works by converting light into an electrical signal. A device such as an oscilloscope
triggers on a certain voltage met when a threshold of light signifying combustion strikes
the diode. The drawback to photodiodes is that only a few set positions are monitored so
only average speed data may be inferred during analysis.
Instead of photodiodes, pressure transducers can be used to help locate the
passing flame front. Heat from the flame and gaseous products increases the pressure as
the flame progresses. In 2007, Bockmon et al, designed an experimental setup that used
both optical and pressure information by combining photodiodes and pressure
transducers. In their experiments, flame speed exhibited a similar trend with pressure
history (Bockmon et al. 2005).
Here, high speed video imaging was used for the flame speed
measurements in an open tray. Loose powders were loaded into an open tray and ignited.
The ensuing videos were analyzed and calibrated. Flame speed was determined by
combining the distance information with the frame rate as a timing device. Drawbacks to
iv
this method, aside from the camera cost, include spatial effects if the camera is not
perfectly perpendicular to the sample. Also, with an optical camera, the luminescent front
is assumed to travel with the flame at the same velocity. For highly convective flows, the
hot product gases may be pushed ahead of the flame at a temperature capable of radiating
visible light. The hot gases may increase the luminescence to a point where
oversaturation leads to systemic error in flame speed calculations.
Conclusion
Flame speeds of micron aluminum particles have been widely studied. With the
onset of nanometer scale aluminum particles, a new regime in solid combustion has
developed. Flame speeds can be 100 times faster just by reducing the size of the fuel
particle (J. Granier and Pantoya 2004). Many theories have been developed in an attempt
to explain this phenomenon including melt dispersion (Levitas et al. 2006), shrinking
core (Park et al. 2005), and reactive sintering (Sullivan et al. 2012). However, none of
these competing theories can completely explain the reaction kinetics for nanoparticle
combustion. This work examines flame speeds in relation to basic reaction parameters in
order to develop an understanding as to which reaction parameters are most important.
Aluminum interactions with halogen containing oxides have only been
studied recently. Flame speeds of 2000m/s have been recorded (Martirosyan, Wang, and
Luss 2009) and product gases (I2, Cl2, F2) have potential in biocidal applications
(Sullivan et al. 2011; B. R. Clark and Pantoya 2010). Zhang et al showed that
mechanically allowing crystalline iodine to Al increased the reactivity of the powder (S.
Zhang et al. 2012). Aluminum particles are passivated with a shell of aluminum oxide
(Al2O3 ). The Al2O3 shell has been combined with iodine to catalyze various reactions in
organic chemistry and to oxidize methane into methanol (Gang et al. 2004). Flourine has
been shown to react with the Al2O3 shell before Al ignites (D.T. Osborne 2006). The
reactions of halogens with the Al2O3 shell leads to the hypothesis that the increased
reactivity is due to interactions between the halogen and the alumina shell. An objective
of this study is to examine the halogen-Al2O3 interaction in order to assess the effect on
flame speed.
iv
CHAPTER II: REACTION KINETICS OF NANOMETRIC
ALUMINUM AND IODINE PENTOXIDE
Introduction
The increase of organized terrorist cells around the world poses a growing
threat to the United States and many other countries. For these terrorist cells, chemical
and biological weapons make highly effective terror weapons against civilians and
weapons of intimidation against soldiers (Harigel 2000). While large scale chemical
weapon production requires a large chemical plant, biological weapons can be produced
in basements and hospitals around the world (Harigel 2000). Of the organisms that could
cause enough disease and death to cripple a region, anthrax poses one of the greatest
threats (Inglesby 1999). Bioweapon attacks from agents such as anthrax would be
difficult to predict, detect or prevent (Inglesby 1999). Therefore, complete elimination of
the bacterial spore while in a storage bunker can effectively prevent great loss of life and
psychological trauma induced from undergoing a terror attack. Popular methods for the
destruction of spore forming bacteria such as anthrax involve either ultraviolet radiation
(Blatchley et al. 2005) or an oxidation agent such as peroxide (Inglesby 1999). An assault
on a bunker storing anthrax containers does not lend itself to a prolonged ultraviolet
radiation exposure. Oxidation of anthrax spores is a slow process with necessary
exposure times of up to an hour for effective neutralization (Blatchley et al. 2005).
Thermites consist of a mixture of Al and a metal oxide which produces a
highly exothermic reaction when ignited (Babaitsev, Arzhevitov, and Presnakova 2007).
With flame temperatures over 2000 K, thermites may act as a quick, effective
sterilization tool when prolonged exposure to a neutralization agent is not a viable option.
A common oxidizer is iron (III) oxide, Fe2O3 (Fan et al. 2006); however, for bacterial
sterilization I2O5 was selected due to iodine’s bactericidal properties (Taylor, Fina, and
Lambert 1970). Iodine pentoxide is produced by heating iodic acid to 200 oC in a stream
of dry air (Holleman and Wiberg 2001). The remaining powder is stable and produces a
thermite reaction when combined with Al fuel. When I2O5 and Al react, the products are
alumina (Al2O3) and iodine.
10𝐴𝑙 + 3𝐼2 𝑂5 → 5𝐴𝑙2 𝑂3 + 3𝐼2
iv
(31)
This study examines the kinetics for nanometer and micron scale Al particles
reacting with I2O5 in a thermal equilibrium experiment from 25 to 1000 oC and in an inert
argon (Ar) environment.
Experimental Method
Four powder samples were prepared consisting of nanometer scale I2O5, 80 nm Al, 15µm
Al, and 40 nm Al2O3. The powder characteristics provided by the suppliers are listed in
Table 1.
Table 1. Powder characteristics
Material Particle Size
Manufacturer
Purity (%)
Al
15µm
Alfa Aesar
99
Al
80nm
Nanotechnologies
80
Al2O3
40nm
Nanotechnologies
99.9
I2O5
Sigma Aldrich
99
Preliminary analysis showed that all the Al is not consumed at the slow heating
rates (10oC/min) investigated. In order to resolve the reaction kinetics, the Al/ I2O5
powders were mixed at varying equivalence ratios () from 0.4 to 1.2 as is defined by Eq.
(32),
𝑚
( 𝑓⁄𝑚𝑜 )
𝑎𝑐𝑡𝑢𝑎𝑙
= 𝑚
(32)
𝑓
( ⁄𝑚𝑜 )
𝑠𝑡𝑜𝑖𝑐ℎ𝑒𝑚𝑒𝑡𝑟𝑖𝑐
where mf and mo are the mass of the fuel and oxidizer, respectively. The Al
particles have an Al2O3 passivation shell accounting for the material’s impurity listed in
Table 1. The Al2O3/I2O5 powder was mixed at a 1:3.27 mass ratio (1:1 molar).
Samples for each of the powders were prepared by suspending the
powders in 60cc of hexane and sonicating the mixture with a Misonix model S3000 for
70 seconds. In order to prevent damaging the oxide shell passivating the Al particles, the
sonicator was programmed to cyclically mix for 10 seconds, and then to stop, allowing
the mixture to cool for 10 seconds. The solutions were then placed in a glass tray under a
fume hood to allow the hexane to evaporate. The powder mixtures were then reclaimed
for further experimentation.
iv
The thermal decomposition of each sample was studied with a Netzsch
STA 409 differential scanning calorimeter and thermo-gravimetric analyzer (DSC/TGA).
The system was programmed to heat the samples at a rate of 10 oCmin-1 from room
temperature to 1000 oC. Sample masses of 7 mg were loaded into the sample crucible,
and the DSC column was evacuated to less than 0.01 Pa using a Pfeiffer model TMU 071
P turbo molecular drag pump. The column was then backfilled with an argon atmosphere
before a 50 mLmin-1 flow of argon was applied to the furnace for the rest of the heating
cycle.
Results and Discussion
The results of the DSC/TG analysis for pure I2O5 can be found in Figure 3.
Figure 3. Heat flow and mass loss curves of I2O5; experiments performed in an Ar
environment at a heating rate of 10°C/min.
On the left axis, the TG % represents the percent mass change occurring at a
given temperature, and on the right , the DSC scale represents the energy change in
mW/mg. Common examples of weight loss are the release of bonded gases, drying of the
sample, and a chemical reaction producing gases. Iodine pentoxide bonds with water to
become iodic acid, HIO3 (Holleman and Wiberg 2001). This reaction becomes reversible
and the acid decomposes at 200 oC as shown by
iv
𝐻𝐼𝑂3 → 𝐻𝐼𝑂3 ∙ 𝐼2 𝑂5 → 1
1
𝐼𝑂
2 2 5
(33)
This decomposition is illustrated in Figure 3 where the bonded water is released
causing the slight mass loss and related endotherm. The I2O5 decomposition was
observed at an onset temperature of 390 °C where the I2O5 fully decomposes into gas
corresponding to a 100% mass loss. It is noted that I2O5 melting and decomposition are
simultaneous (Holleman and Wiberg 2001).
1
𝐼2 𝑂5 → 𝐼2 + 2 𝑂2
2
(34)
The I2O5/ Al2O3 reaction is displayed in Figure 4.
Figure 4. Heat flow and mass loss curves for Al2O3 and I2O5, experiments performed in
an Ar environment at a heating rate of 10°C/min.
By combining I2O5 and Al2O3, the kinetics between I2O5 and the Al2O3 shell
encapsulating the Al particles were isolated. The main differences between Figure 3 and
Figure 4 are the presence of an inflection in the mass loss curve and the presence of a
third endotherm with an onset temperature of 460 °C. The inflection point in the mass
loss (Figure 4) may correspond to a slight exothermic behavior causing the heat flow
curve to return to the baseline slighter faster than in Figure 3. The inflection and
iv
exotherm may be described by Al2O3 binding to either the I2 or O2 gas when these oxides
disassociate. Since Al2O3 is inert to oxygen at this temperature, the reaction must be due
to the Al2O3 binding to the iodine when freed from disassociating I2O5. This is an
interesting observation and similar to a study by Osborne and Pantoya (Dustin T.
Osborne and Pantoya 2007) that showed fluoride ions from Teflon decomposition bond
to the alumina shell passivating Al particles. Sarbak (Sarbak 1997) showed the
interaction between fluoride ions and Al2O3 may be facilitated by hydroxyls that are
bonded to portion of the Al2O3 surface. Both of these studies showed the fluoride ion
sorption into alumina was an exothermic reaction (Dustin T. Osborne and Pantoya 2007;
Sarbak 1997). Toyohara et al (Toyohara et al. 2002) studied the iodine sorption
mechanism into mixed solid alumina cement. They similarly found iodine ions replaced
hydroxyls and bonded to alumina. Fluorine and iodine are halogens and this study (as
well as others (Dustin T. Osborne and Pantoya 2007; Sarbak 1997; Toyohara et al. 2002))
suggests both have similar sorption behaviors when interacting with alumina.
At 550oC mass loss in Figure 4 continues again accompanied by a third
endotherm. This second stage of mass loss and related endotherm may correspond to
Al2O3 phase changes from amorphous Al2O3 to - Al2O3 (M Trunov et al. 2005). The
alumina phase change may trigger a release of iodine gas resulting in a subsequent mass
loss.
Figure 5 shows the analysis of the reaction between micron scale Al and I2O5.
iv
Figure 5. Heat flow and mass loss curves for micron scale Al and I2O5, experiments
performed in an Ar environment at a heating rate of 10°C/min. Note the lack of any
exothermic reaction
iv
Passivated micron Al has been shown to be thermally nonreactive until it
approached its melting temperature of 660oC (J. J. Granier and Pantoya 2004). Also, the
lower specific surface area of the larger particles reduces the Al2O3 shell/I2O5 contact
area limiting the sorption of iodine into the amorphous Al2O3. This specific surface area
affect was also observed for the sorption of fluoride ions on varying specific surface area
Al particles (Dikici et al. 2009). A combination of the lower specific surface area of the
Al particles coupled with the 400 oC degradation of I2O5 and the slow heating rate gives
the argon gas flowing through the cylinder 22 minutes to purge the oxygen and iodine
gases from the system resulting in the lack of any reaction.
Nanometer scale Al, however, has been shown to be reactive in its solid state at
much lower temperatures (Dustin T. Osborne and Pantoya 2007; J. J. Granier and
Pantoya 2004). This increased reactivity allows the Al to react with the decomposed I2O5
as shown in Figure 6 resulting in a complex multistep process.
Figure 6. Nanometer scale Al powder and I2O5 with an equivalence ratio of 0.4,
experiments performed in an Ar environment at a heating rate of 10°C/min.
iv
Similar to Figure 3, an endotherm appears at 200 oC resulting from iodic acid
releasing water and forming I2O5. Also, the endotherm at 400 oC corresponding to the
decomposition of I2O5 can be observed but is somewhat masked by the exothermic
reaction of Al and oxygen with an onset temperature of 309 °C shown in Eq (35).
(35)
4𝐴𝑙 + 3𝑂2 → 2𝐴𝑙2 𝑂3
Examination of the TG curve reveals a similar region to Figure 4 over 400 to 550
o
C implying an interaction between the Al2O3 passivation shell and the iodine gas.
Examination of the DSC curve over this temperature range reveals that the endotherm for
the iodine release at 400 °C to the alumina phase change at 550 °C is partially masked by
an exothermic Al-I-O reaction.2𝐴𝑙 + 3𝐼2 → 2𝐴𝑙𝐼3
Finally, an endotherm can be seen at 660oC where unreacted Al melts.
iv
CHAPTER III: REACTION KINETICS OF NANO-ENERGETICS
WITH HALOGEN CONTAINING OXIDES
Introduction
Halogenated oxidizers have received significant attention recently for their
potential application in neutralizing spore forming bacteria via the production of halogen
gases (B. R. Clark and Pantoya 2010; Sullivan et al. 2011; S. Zhang et al. 2012). Due to
the electronegativity of common halogens: fluorine (4.0), chlorine (3.0), iodine (2.5) and
oxygen (3.5) (Brown, Lemay, and Bursten 2000), reactions containing oxygen with either
iodine or chlorine have the potential for the aluminum to react with oxygen leaving the
halogen species as a product. Fluorine, while potentially effective for neutralizing
bacterial spores, is simply too reactive to be utilized in this application because it bonds
too quickly with surrounding gasses or aluminum leaving the spore unaffected.
Therefore, iodine and chlorine containing oxides were selected for this analyses including
I2O5, AgIO3, KClO4, Ca(IO3)2 each combined with aluminum (Al) fuel. The goal is to
analyze the relationship between reaction parameters for halogen containing formulations
and advance our understanding of nano-energetic material composite reactivity.
It has been well documented that energetic materials undergo a transition in flame
speed as particle size decreases (J. Granier and Pantoya 2004; Lynch, Krier, and Glumac
2009; Moore, Son, and Asay 2004; Perry et al. 2004; Prentice, Pantoya, and Clapsaddle
2005; Tillotson et al. 2001). With materials on a nanometer scale, flame speeds have been
documented on the order of 2000m/s for Al+I2O5 (Martirosyan, Wang, and Luss 2009).
While many theories exist as to why flame speeds are so high with certain formulations, a
commonly accepted hypothesis is that mixtures such as Al+MoO3 and Al+I2O5 produce
high flame speeds because they produce significant heat and gas that propel energy
forward in a pressure wave convecting rapidly to the yet unburned reactants (Asay et al.
2004). However, in an unconfined or open-tray burn, convective effects are reduced
without any confinement directing pressure to move in the direction of the unburned
powder. For example, in this study, the flame speed for Al+I2O5 in an unconfined burn
was found to be 916m/s as opposed to confined flame speed measurements of 2000 m/s
made by Martirosyan (Martirosyan, Wang, and Luss 2009). In addition to the role of
iv
confinement, there is a need for more research investigating reaction parameters and their
effect on flame speed.
When considering combustion of solid powders from a fundamental perspective,
an energy threshold, defined here as apparent activation energy, must be achieved in
order for a reaction to initiate. Once ignited and in order for the reaction to selfpropagate, the heat generated by the reaction, called heat of combustion, must be
sufficient in that enough energy is transferred to the adjacent particles to elevate them to
a state beyond their activation energy. In 2006, Trunov et al. showed that the smaller
particle size resulted in increased reactivity as well as a reduced melting point in Al
particles (MA Trunov 2006). As particle size is reduced from micrometer to nanometer
scales, the apparent activation energy decreases to the point of igniting the aluminum
reaction even in the solid phase at temperatures well below Al melting point at 660oC
(Sanders et al. 2007). Since the transition from micron to nanoscale particles results in
both an increase in flame speed as well as a decrease in apparent activation energy, one
could hypothesize a correlation between the two. Also, these macroscopic combustion
behaviors are linked to molecular reactions at the interfaces between grain boundaries
within the system. Linking the macroscopic observations with molecular level kinetics
will enable a more complete understanding of the overall reaction dynamics. Therefore,
an objective of this study is to examine this correlation by experimentally comparing
flame speed to apparent activation energy, gas production, and heat of combustion and
further modeling molecular dynamic interactions between grain boundaries to resolve
fundamental kinetics governing the overall reaction mechanism. This molecular dynamic
analysis is limited to the kinetics of the fastest burning mixtures. The combined
experimental/theoretical picture presented in this work shows the complexity of Al –
halogen oxide reactions. At the same time, the results reveal some hallmark features of
reactions involving Al fuel that includes alumina abstracting oxide species and forming
intermediates that increase the exothermic yield.
Methods
The Al powder has an 80nm average particle diameter and supplied by
Novacentrix (formerly Nanotechnologies), Austin, TX. Particles are passivated with a 4
iv
nm thick alumina shell such that the Al content is 75 % by volume. For the oxidizers,
I2O5 and KClO4 were supplied by Sigma Aldrich, St Louis MO, and AgIO3 and Ca(IO3)2
by Alfa Aesar, Ward Hill, MA. See Table 2 for powder information.
Table 2. Reactant average particle diameter, supplier and anticipated products from Al
reaction with respective oxidizer
Powder
Diameter
Supplier
Anticipated Products
Al
I2O5
Ca(IO3)2
AgIO3
KClO4
80nm
355µm
15µm
100nm
50µm
Novacentrix
Sigma Aldrich
Alfa Aesar
Alfa Aesar
Sigma Aldrich
Al2O3+I2
Al2O3+CaI2
Al2O3+AgI
Al2O3+KCl
All reactions were balanced based on oxygen transfer assuming 100% of the
oxygen transferred to the aluminum. Samples for each of the powders were prepared by
suspending the powders in 60 cc of hexane and sonicating the mixture with a Misonix
model S3000 for 70s. In order to prevent damaging the alumina shell passivating the Al
particles, the sonicator was programmed to cyclically mix in intervals of 10s allowing the
mixture to cool for 10s. The solutions were then placed in a glass tray under a fume hood
to allow the hexane to evaporate. The powder samples were then reclaimed for further
experimentation. In all experiments, loose powder was used, such that the powder
samples are 5% of the theoretical maximum density (TMD).
Near Equilibrium Experiments
Both thermal equilibrium and non-equilibrium tests were performed and the
results compared to thermal equilibrium models (i.e. REAL code modeling software) in
order to gauge how the model predicts actual energy output and relative efficiencies of
each reaction. A Neztsch STA 409 differential scanning calorimeter and
thermogravimetric analyzer (DSC/TGA) was used for thermal equilibrium testing. An
8mg sample was loaded into a platinum crucible with an alumina liner and placed on a
sample carrier inside an oven and heated at rates of 2, 5, and 10K/min until the reaction
was complete. The crucible lid is vented in order to prevent pressurization within the
iv
device. Within the DSC/TGA, the sample crucible is compared to an empty reference
crucible in order to obtain the net energy and mass changes. Also, the sample carrier is
mounted on a microscale (i.e. TGA) allowing for mass change measurements that relay
phase change (i.e. gas production) information as a function of equilibrium temperature.
The slope of the DSC curve changes when the reaction within the DSC/TGA produces
enough energy to become noticeable within the natural noise of the machine. The
temperature at which the slope changes is defined here as the onset temperature. The area
under the DSC curve corresponds with the net exothermic behavior. Apparent activation
energy (Ea) can be determined from the DSC measurements using the Ahhrenius
equation in Eq. (1) that is described as the Type B-1.95 Peak Method described by M.
Starink (Starink 2004), the procedure for which can be found in ASTM 698E (ASTM
E698 2011).
In this method, the reaction rate is estimated as the ratio of heating rate (B) to
temperature at which the reaction is at its peak (Tp) and R is the universal gas constant.
The pre-exponential factor, A, is a constant also expressed as C in Eq. (2).
𝐵
𝐸𝑎
)
(36)
𝑅𝑇𝑝
By taking the natural log of both sides, 𝐸𝑎 can be calculated by plotting the
𝑇𝑝1.95
ln (
𝑇𝑝1.95
𝐵
= 𝐴 exp (−
1
) vs 𝑅𝑇 for each of the heating rates and finding the slope of the best fit line.
𝑝
𝑇𝑝1.95
𝐸𝑎
ln (
)=
+𝐶
𝐵
𝑅𝑇𝑝
(37)
Non-Equilibrium Experiments
For the non-equilibrium tests, a Parr model 1341 bomb calorimeter was used with
a model 1108 oxygen combustion bomb chamber as seen in Figure 7. Using REAL code
(Timtec L.L.C.) values for theoretical heats of combustion based on the heat of formation
for the reactants and products, powder mass was determined in order to obtain a 0.5 to 1K
elevation of a 2kg water bath above the calibrated energy inputs of a stirrer and other
environmental factors. A thermocouple within the water bath was sampled at a rate of 10
Hz for 40 minutes. Using a 60% value between the lines extrapolated before and after
iv
the reaction on a temperature as a function of time plot, a change in temperature was
obtained in order to find the energy associated with heating the steel bomb chamber and
water.
Figure 7. Schematic of bomb calorimetry experiment.
For the flame speed experiments, an average of 500 mg of loose powder was
loaded into a 0.5 by 13cm notch and distributed evenly with no packing. By taking the
volume of the notch and the mass of powder added, %TMD was calculated as an average
of 4.5%. Nickel-chromium wire, Ni80-020, from Omega Engineering, Inc was placed at
one end of the notch and heated at 7V by a variac transformer. Once the powder ignited,
flame speeds were recorded via a Phantom 7 high speed camera at frame rates ranging
from 900 to 90000 frames per second as seen in Figure 8. National Instruments Vision
Builder AI 2010 software was used to locate the flame front in the high speed images and
when combined with the frame rate, allows for the calculation of flame speed.
Figure 8. Schematic for open tray flame speed measurements.
iv
iv
Ab Initio Quantum Chemical Calculations
All simulations were performed for the Al – I2O5 system. This reaction can be
described by chemical events at the Al grain boundaries. The role of grain boundaries in
initiation of chemical reactions is complex. However, if the events are broken down in
solid-state and gas-phase reactions, the state of initial and final phases of reactants are
important aspects of the process. In order to use first-principles calculated
thermokinetics to describe the complex series of elementary reactions, the following
assumptions are applied. First, Al particles are covered with an Al2O3 oxide layer. This
oxide layer is often amorphous but modeling is performed using stable -Al2O3 as the
ground state surface. In general, the first-principles calculated reactions show the upper
limit for kinetic barriers for such reactions. Rate observed is higher as Al2O3 layers
containing defects and disorders are more reactive than ideal surface morphology used in
the current work.(Wang et al. 1999; Zhai and Chen 2011; L. Zhang et al. 1995) Thermal
excitations at the Al/Al2O3 grain boundaries in presence of oxidizers are a primary focus
when studying aluminized reactive composites. The working hypothesis of the
simulations presented in this work assumes that upon activation of reactions at the
Al/Al2O3/I2O5 interfaces, exothermic chemical reactions simultaneously weaken the
oxide layer and increase local temperatures. This mechanothermal load present during the
initiation stages can lead to melting of core Al, rupture of oxide layer, acceleration due to
faster Al-I2O5 reactions and related species, and eventual ignition or detonation
depending on the reaction conditions.
An objective of this analyses is to explore the Al2O3-I2O5 interfaces using firstprinciples calculated initial adsorption (exothermic) and possible I-transfer (activated
endothermic) processes at the interface using condensed-phase periodic density
functional theory (DFT) calculations. The gas-phase thermokinetics of the reactions are
determined by three primary groups of coupled reactions: (a) Al released undergoes
exothermic oxidation in air producing Al2O3 (b) I2O5 thermally decomposes at
temperatures produced by oxidation reactions and (c) Al reacts with I2O5 decomposition
fragments. Such reactions are studied for their reaction enthalpies using gas-phase
quantum chemical theories at various levels including spin-orbit DFT (SO-DFT) level
theories needed for I2O5 decomposition reactions.
iv
In the present work we describe DFT and SO-DFT quantum chemical calculations
on a selection of iodine oxide molecules and their reactions with aluminum which are
important initial steps in the mechanistic understanding of Al/I2O5 thermite reactions.
First, the elementary steps and the associated thermochemistry for the dissociation of the
I2O2, I2O3, I2O4, and I2O5 oxides. Then , reactions between the [Al13]- cluster and iodine
oxide dissociation fragments will be discussed. Our recent theoretical study (Losada and
Chaudhuri 2010) of the finite size effects on aluminum / polytetrafluoroethylene (PTFE)
reactions showed that the [Al13]- cluster provides a suitable model for some aspects of the
surface of bulk Al or nanoscale Al in reactive composites. This system is the smallest one
in which an Al atom can have a coordination of 12 of a close-packed bulk crystal and the
surface atoms mimic the triangular Al arrangement on an Al(111) surface termination. In
addition, [Al13]- has been proved, both experimentally and theoretically, (Hettich 1989;
Rao and Jena 1999) to be a stable system. Its stability arises from closures of its geometry
and electronic shells. With heavy element iodine present in determining the chemical
reactivity in Al/I2O5 composites, it is necessary to include relativistic effects when
calculating the thermochemistry of these systems.
Solid State Calculations
An -Al2O3 layer oriented in (100) direction was used for studying I2O5
adsorption. An ultrasoft pseudopotential and revised Perdew Burke Ernzerhof (RPBE)
exchange correlation functional as implemented in DFT code CASTEP was used for all
calculations. (S. J. Clark et al. 2005; Segall et al. 2002) The slab thickness varied until
convergence was achieved at kinetic energy cut-off of (300 eV) and a surface layer
consisted of 160 atoms in a 9.83 x 9.83 Å surface in an approximately 16 Å thick slab.
Due to this large size of the unit cell needed to accommodate the precision in surface
energy and relatively large I2O5 molecule, -point calculations were sufficient to obtain
convergence with respect to known surface energy values for Al2O3. An I2O5 molecule
was first optimized using gas-phase DFT calculations (described below). The optimized
I2O5 molecule was allowed to relax in a same size period box for RPBE calculated total
energy for a gas-phase molecule (EIO). The I2O5 molecule was allowed to adsorb on
Al2O3 (100). The following energy difference was used to estimate absorption energy:
iv
Eads=ESurface+EIO –ESurface+IO. The I2O5 adsorbed on Al2O3 surface using a strong I-O-Al
bridge bond.

Gas-phase Calculations.
In selecting a suitable and effective methodology for the treatment of the
molecular systems of I2O5 and fragments considered, which contain the heavy iodine
atom, oxygen and aluminum, the choice of appropriate basis set is crucial for the
accuracy of the calculations. Thus, correlation consistent basis sets were used. More
specifically, the diffuse-augmented sets from double- to quadruple-zeta were used for
oxygen, aug-cc-pVnZ (n = D, T, Q), (Dunning 1989; Kendall, Jr, and Harrison 1992)
while those based on accurate relativistic small-core pseudopotentials/correlation
consistent basis sets developed by Peterson et al.(Peterson et al. 2006) were taken for
iodine, aug-cc-pVnZ-PP (n = D, T, Q). In the following, the combination of the aug-ccpVnZ on oxygen and aug-cc-pVnZ-PP on iodine is abbreviated as aVnZ.
The B3LYP, (C. Lee, Yang, and Parr 1988; Becke 1993) PBE0, (Adamo and
Barone 1999) ACM, (Becke 1993) M05, (Zhao and Truhlar 2004) and M06-L (Zhao and
Truhlar 2006) functionals were used. Computational studies of the enthalpies of
formation of OXO(S. Y. Lee 2004) (X = Cl, Br, and I) showed that the B3LYP functional
delivers good results when compared with CCSD(T) correlated calculations. As for the
PBE0 and ACM functionals, the investigation of Lee and co-workers showed that these
functionals performed reasonably well for molecular geometries and vibrational
frequencies in iodine-containing molecules.(Kim, Ihee, and Lee 2010) Moreover, the
performance of M05 and M06-L for iodine oxides has not been assessed for these oxide
molecules. For these reasons, we chose these five DFT functionals for the present study.
Structure optimizations and harmonic frequencies were calculated at the corresponding
levels for the aVDZ and aVTZ basis. Zero-point energy (ZPE) and enthalpy corrections
for the aVQZ level were evaluated using the harmonic frequencies calculated at the
aVTZ level for each functional.
In addition to the scalar relativistic effects which are taken into account by using
the relativistic small-core pseudopotential/correlation consistent basis sets, a second
iv
relativistic correction, spin-orbit (SO), was applied to all the iodine oxide molecules in
this work. The effects due to SO were recovered by use of the SO-DFT method as
implemented in the NWChem program.(Valiev et al. 2010) SO effects were calculated
using the aug-cc-pVnZ-PP SO (n = D, T, Q) basis set on iodine and the corresponding
aug-cc-pVnZ on oxygen for each DFT functional considered. We used uncontracted basis
for the SO-DFT calculations to get a better description of the SO effects, as pointed out
by Lee and co-workers.(Kim, Ihee, and Lee 2010)
Heats of formation were calculated by combining our computed atomization
energy values with the known enthalpies of formation (fH) at 0 K for oxygen and
iodine, fH (O) = 58.98 kcal/mol and fH (I) = 25.61 kcal/mol.(Grant et al. 2010)
Enthalpies of formation at 298 K were obtained by following the procedures outlined by
Curtis et al.(Curtiss et al. 1997)
For the calculations of the aluminum-iodine oxide polyatomics, we used the
[Al13]- cluster as the model system. Geometries were optimized using the all-electron
aug-cc-pVTZ basis for oxygen and the aug-cc-pVTZ-PP for iodine in conjunction with
the 6-311G basis set for aluminum. This level of theory was previously used for the
description of the [Al13]- -Teflon combustion reactions.(Losada and Chaudhuri 2010)
Because of the computational expense of geometry optimization and frequency
calculations with aug-cc-pVQZ basis sets for these large systems, the aug-cc-pVTZ
geometries were used in single point aug-cc-pVQZ calculations. Harmonic frequencies
were calculated using the aug-cc-pVTZ on oxygen, aug-cc-pVTZ-PP on iodine, and 6311G basis set on aluminum. No symmetry constraints were imposed in the optimization
of all the molecular species. All calculations for the aluminum-iodine oxide polyatomics
were carried out with the Gaussian 09 suite of programs and the SO-DFT calculations
with the NWChem program.
iv
Results
Figure 9. Heat flow measurements for Al+I2O5 scanned at 2, 5 and 10 K/min heating
rates.
𝑇𝑝1.95
Figure 10 shows ln (
𝐵
1
) as a function of 𝑅𝑇 , the resulting slope is Ea .
𝑝
Figure 10. Peak temperatures for Al+I2O5 plotted using the Type B-1.95 Peak Method.
The resulting Ea is 175.6 kJ/mol
iv
The activation energy can be converted to a mass basis since applied energy is
commonly considered in the same physical sense. For this reason, Ea for each material is
converted to kJ/g in . Uncertainty for this method in calculating Ea is largely based on
repeatability in finding the peak temperatures in the DSC, and for our device, Ea values
have been within 15%.
The results for a typical bomb calorimetry experiment are shown in Figure 11.
Figure 11. Graph showing Al+I2O5 Parr bomb calorimetry result displaying dT
In order to calculate the heat of combustion, Hc = CpdT was used where Hc is the
heat of combustion, Cp is the specific heat of the water bath, and dT is the change in
temperature as seen in Figure 5. The specific heat was modeled to include a mass
percentage of the specific heats of the steel container (3.838 kg, 473 J/kgK) and the water
bath (2.00kg, 4184J/kgK) for a final value of 10.18kJ/K. Also, Hc was found for each
mixture and was found to be 3.68±0.22 kJ/g for Al+I2O5, 2.64±0.20 for Al+Ca(IO3)2,
2.67±0.30 for Al+AgIO3, and 3.62±0.16 for Al+KClO4. The average values are shown in
Table 3.
Representative still frame images of the flame propagating are shown in Figure
12.
iv
Figure 12. Open tray still frame images A) Al+I2O5 taken at 33µs intervals, B)
Al+Ca(IO3)2 at 17ms intervals, C) Al+AgIO3 at 40 µs intervals, D) Al+Fe2O3 at 54ms
intervals.
The resulting analysis from the luminescent front signaling flame propagation is
displayed below in Figure 13.
Figure 13. Position of flame front as a function of time for one test for Al+I2O5 with a
flame speed of 898 m/s.
For the fast class of powders, Al+I2O5 and Al+AgIO3, the average flame speed
was found to be 916±53 and 926±29 m/s, respectively. For the slow class of powders,
iv
Al+KClO4 and Al+Ca(IO3)2, the average flame speed was found to be 0.77±0.09 and
2.50±0.92 m/s, respectively. The results for the average flame speed for each mixture
have been tabulated in Table 3.
Table 3 displays the compilation for the experimental results over the three
different testing protocols for the four reaction chemistries examined. The heat of
combustion from a thermal equilibrium code (REAL), an equilibrium diagnostic (DSC),
and a non-equilibrium method (Bomb) are displayed as well as mass loss from the TGA,
apparent activation energy (Ea), onset temperature, and average flame speed.
Table 3. Summary of experimental results.
Reactants
10Al+3I2O5
4Al+Ca(IO3)2
2Al+AgIO3
8Al+3KClO4
TGA Mass
Loss (%)
95
75
35
0
Heat of Combustion,
(kJ/g)
REAL
6.14
7.81
5.02
10.63
DSC Bomb
1.13 3.68
0.18 2.64
1.49 2.67
2.31 3.62
Ea
(kJ/g)
Onset
Temperature
(oC)
Average
Speed
(m/s)
1.79
2.78
1.84
3.62
310
511
309
514
916
2.5
926
0.8
The results from theoretical calculations are summarized in Table 4 and Table 5.
The thermal decomposition enthalpies for different iodine clusters calculated at B3LYP
and PBE0 level are provided in Table 4. The reactions are all exothermic when O2 is
produced from an iodine oxide cluster. So, even if iodine oxide’s role is limited to
producing I2 and O2, Table 5 shows that such steps also add to the highly exothermic
nature of the Al/I2O5 reaction.
iv
Table 4. Decomposition of iodine oxide clusters in gas phase for B3LYP, PBE0 and
ACM functionals at aug-cc-pVTZ (aVTZ) and aug-cc-pVQZ (aVQZ) levels of theory.
The values in the bracket are calculated using spin-orbit (SO) corrections.
B3LYP
ACM
r H
r G
r H
r G
rH
r G
37.8 [38.5]
39.2 [36.4]
25.3 [26.3]
26.7 [24.2]
42.6 [43.9]
44.4 [41.2]
29.8 [31.4]
31.6 [28.7]
39.9 [40.8]
41.5 [35.9]
27.7 [28.8]
29.3 [23.7]
aVTZ
aVQZ
-12.0
[-11.0]
-8.8
[-12.2]
-22.3
[-21.9]
-19.1
[-23.3]
-9.8 [-8.9]
-6.6 [-10.5]
-20.0
[-19.8]
-16.8
[-21.5]
-9.0 [-8.3]
-6.2 [-9.6]
-18.7
[-19.1]
-15.8
[-20.5]
aVTZ
aVQZ
6.1 [6.7]
6.6 [7.3]
-5.4 [-5.0]
-4.9 [-4.4]
9.5 [10.5]
10.1 [10.9]
-2.1 [-1.3]
-1.5 [-0.9]
6.5 [7.8]
7.3 [8.4]
-5.0 [-3.6]
-4.1 [-3.0]
-11.2 [-9.8]
-8.6 [-8.2]
-21.3
[-19.9]
-18.7
[-18.3]
-12.4
[-10.4]
-9.5 [-8.8]
-22.0
[-20.6]
-19.7
[-19.0]
-11.7 [-9.8]
-9.3 [-8.4]
-22.0
[-20.1]
-19.6
[-18.7]
35.4 [34.8]
36.2 [35.3]
25.3 [24.7]
26.1 [25.2]
38.7 [38.8]
39.7 [39.4]
28.6 [28.7]
29.5 [29.3]
37.1 [37.0]
37.9 [34.7]
27.0 [26.8]
27.8 [24.6]
23.2 [22.3]
24.5 [22.9]
12.1 [11.2]
13.4 [11.9]
27.3 [27.1]
28.9 [27.6]
16.2 [16.0]
17.3 [16.5]
24.6 [24.3]
26.4 [25.0]
13.5 [13.2]
15.3 [14.0]
-6.9 [-6.7]
-3.4 [-5.2]
-17.9
[-17.6]
-14.3
[-16.7]
-5.3 [-4.7]
-1.4 [-3.0]
-16.3
[-15.7]
-12.4
[-14.0]
-4.8 [-4.3]
-1.0 [-2.7]
-15.8
[-15.0]
-12.0
[-13.7]
20.0 [17.5]
21.4 [17.9]
11.0 [8.5]
12.4 [8.9]
25.7 [23.4]
27.1 [24.1]
16.7 [14.4]
18.1 [15.1]
24.4 [22.0]
25.8 [22.6]
15.4 [13.0]
16.8 [13.6]
17.0 [11.5]
17.7 [11.6]
8.5 [3.0]
9.2 [3.2]
20.8 [16.1]
21.4 [16.4]
12.3 [7.6]
12.9 [8.0]
19.3 [14.7]
20.3 [15.0]
10.8 [6.2]
11.8 [6.5]
-25.8
[-24.3]
-22.9
[-23.3]
-34.8
[-33.3]
-32.0
[-32.2]
-26.2
[-24.7]
-23.2
[-23.3]
-35.3
[-33.7]
-32.3
[-32.4]
-24.5
[-22.9]
-21.8
[-21.7]
-33.6
[-32.0]
-30.9
[-30.7]
Decomposition
Basis Set
I2O5 → IO2 + IO3
aVTZ
aVQZ
→ I2O3 + O2
I2O4 → IO2 + IO2
PBE0
→ I2O2 + O2
aVTZ
aVQZ
→ IO + IO3
aVTZ
aVQZ
I2O3 → IO + IO2
aVTZ
aVQZ
→ I2O + O2
aVTZ
aVQZ
aVTZ
aVQZ
I2O2 → IO + IO
→ IO2 + I
→ I2 + O2
aVTZ
aVQZ
aVTZ
aVQZ
As concentration of gas-phase I2O5 is limited in high temperatures, reactions with
smaller fragments are shown in Table 5.
iv
Table 5. The reaction enthalpy and free energy calculated for Al reactions with iodine
oxide fragments calculated at aug-cc-pVTZ (aVTZ) and aug-cc-pVQZ (aVQZ)
PBE0
Reaction
Basis
Seta
ACM
r H
r G
r H
r G
[Al13]- + I2 → [Al13]- -I + I
(1)
aVTZ
aVQZb
-12.8 [-16.6]
-12.1 [-15.9]
-11.9 [-15.7]
-11.2 [-15.0]
-11.8 [-15.5]
-11.0 [-15.0]
-10.7 [-14.3]
-9.8 [-13.8]
[Al13]- + IO → [Al13]- -I + O
(2)
aVTZ
aVQZ
-1.3 [1.9]
-0.6 [2.8]
-0.70 [2.5]
-0.02 [3.4]
1.8 [5.2]
2.5 [6.0]
2.7 [6.1]
3.4 [6.9]
→ [Al13]- -O + I
(3)
aVTZ
aVQZ
-69.3 [-72.4]
-69.5 [-73.1]
-67.3 [-70.4]
-67.5 [-71.1]
-69.8 [-72.9]
-70.0 [-73.6]
-67.8 [-70.9]
-68.0 [-71.6]
[Al13]- + I2O → [Al13]- -I + IO (4)
aVTZ
aVQZ
-25.9 [-24.3]
-24.8 [-22.7]
-25.4 [-24.0]
-26.5 [-22.4]
-25.0 [-23.3]
-23.9 [-21.8]
-24.4 [-22.7]
-23.3 [-21.2]
→ [Al13]- -O + I2 (5)
aVTZ
aVQZ
-82.3 [-80.0]
-80.5 [-82.8]
-80.8 [-78.6]
-79.0 [-81.4]
-83.0 [-80.8]
-81.2 [-82.4]
-81.6 [-79.4]
-79.8 [-80.9]
→ [Al13]- -IO + I (6)
aVTZ
aVQZ
-35.6 [-39.0]
-37.3 [-39.2]
-33.3 [-36.8]
-35.0 [-37.0]
-35.8 [-38.0]
-36.1 [-38.3]
-34.0 [-36.2]
-34.3 [-36.5]
[Al13]- + IO2 → [Al13]- -I + O2 (7)
aVTZ
aVQZ
-65.2 [-62.7]
-63.5 [-61.8]
-64.8 [-62.4]
-63.2 [-61.5]
-61.7 [-59.1]
-60.1 [-58.5]
-61.0 [-58.5]
-59.5 [-57.9]
→ [Al13]- -O + IO (8)
aVTZ
aVQZ
-68.5 [-69.3]
-67.8 [-70.7]
-67.0 [-67.8]
-66.3 [-69.2]
-69.0 [-69.8]
-68.3 [-71.1]
-67.5 [-68.3]
-66.8 [-69.6]
[Al13]- + IO3→[Al13]- -IO + O2 (9)
aVTZ
aVQZ
-94.3 [-92.1]
-92.6 [-91.4]
-93.8 [-91.6]
-92.0 [-90.9]
-91.5 [-89.4]
-89.8 [-86.0]
-91.6 [-89.5]
-89.7 [-86.1]
→ [Al13]- -O + IO2 (10)
aVTZ
aVQZ
-87.3 [-86.5]
-86.8 [-85.0]
-87.1 [-86.3]
-86.6 [-84.8]
-88.8 [-88.0]
-88.1 [-86.7]
-88.6 [-87.9]
-87.8 [-86.6]
[Al13]- + I2O2→[Al13]- -I +IO2 (11)
aVTZ
aVQZ
-31.7 [-29.8]
-30.4 [-28.0]
-33.9 [-32.0]
-32.6 [-30.2]
-30.8 [-29.0]
-29.3 [-28.1]
-31.4 [-29.6]
-29.8 [-28.7]
→ [Al13]- -O + I2O (12)
aVTZ
aVQZ
-74.4 [-74.9]
-73.5 [-75.2]
-75.4 [-75.9]
-74.5 [-76.2]
-74.7 [-75.5]
-74.0 [-75.8]
-74.4 [-75.1]
-73.6 [-75.4]
→ [Al13]- -IO + IO
(13)
aVTZ
aVQZ
-41.9 [-41.3]
-40.7 [-42.2]
-42.6 [-42.0]
-41.4 [-42.9]
-40.7 [-40.2]
-39.1 [-41.3]
-40.6 [-40.0]
-38.9 [-41.1]
→ [Al13]- -IO2 + I
(14)
aVTZ
aVQZ
-34.7 [-34.2]
-33.3 [-37.6]
-32.5 [-32.0]
-31.1 [-35.4]
-32.9 [-32.1]
-31.5 [-36.5]
-30.8 [-30.0]
-29.4 [-34.4]
[Al13]- +I2O3→[Al13]- -IO+IO2 (15)
aVTZ
-35.2 [-32.8]
-36.0 [-33.3]
-35.6 [-33.0]
-37.0 [-34.4]
iv
aVQZ
-34.8 [-32.3]
-35.3 [-32.8]
-34.9 [-32.7]
-36.3 [-34.1]
→ [Al13]- -IO2 + IO (16)
aVTZ
aVQZ
-27.2 [-26.2]
-26.4 [-25.2]
-26.6 [-25.6]
-26.0 [-24.6]
-26.7 [-23.7]
-25.7 [-25.4]
-26.6 [-23.6]
-25.6 [-25.3]
→ [Al13]- -IO3 + I (17)
aVTZ
aVQZ
-44.7 [-48.6]
-45.1 [-47.8]
-42.0 [-45.9]
-42.4 [-45.1]
-43.2 [-47.3]
-43.7 [-46.6]
-39.7 [-43.4]
-39.8 [-42.7]
aVTZ
aVQZ
-61.8 [-60.3]
-60.0 [-61.7]
-61.5 [-59.9]
-59.6 [-61.4]
-63.8 [-62.2]
-62.7 [-63.0]
-63.9 [-61.1]
-62.8 [-62.8]
→ [Al13]- -O + I2O2
(18)
a aVTZ is for aug-cc-pVTZ on oxygen and aug-cc-pVTZ-PP basis on iodine atoms. Geometry optimization and
harmonic frequency calculations were carried out at the aVTZ level for oxygen and iodine in conjunction with the 6311G basis for aluminum. b aVQZ is for single point calculations at the optimized aVTZ geometries.
Discussion
One of the first points to notice is that the heats of combustion are much lower in
equilibrium experiments than in non-equilibrium and both of these measurements are
significantly lower than predicted by the modeling software. For the equilibrium study
(DSC), the difference can be explained by the reactants or intermediates gasifying and
leaving the system prior to reaction, thus not contributing to the heat production. The off
gassing cannot be calibrated by comparison to a known standard because different
reactions will experience phase changes at different paces. Therefore, non-equilibrium
tests provide a more accurate indication of the actual heat of combustion for a given
reaction.
From Table 3 there is not a direct correlation to flame speed for mass loss or heat
of combustion. By using the thermo-equilibrium TGA experiments, the gas generation
for a reaction is approximated to an order of scale. If a convective flow was actually
driving the unconfined reaction to produce the higher flame speeds, one would expect to
see high mass loss and a large heat of combustion. However, the only factor which
corresponds to a fast burning powder is apparent activation energy.
If activation energy as defined by the Arrhenius equation was the primary
controller for flame speed, an exponential relationship would be expected between the
reaction rate and Ea. For the four mixtures chosen for this study, changes in Ea result in
order of magnitude changes in flame speed. Other factors may be contributing to this
relationship, but the dramatic changes in flame speed correlating to small changes in Ea
iv
would imply that the chemical kinetics of the reaction are the primary drivers for flame
speed in open tray burns instead of a physical limiting agent such as diffusion rates, mass
transfer, etc. For this reason, the remainder of this section is focused on the intermediate
chemistry influencing the high flame speeds observed in Table 3.
Halogens have been shown to both bond to alumina (Farley and Pantoya
2010; Saikia et al. 2009; Watson, Pantoya, and Levitas 2008) and accelerate aluminum
oxidation (Gang et al. 2004). For example, Osborne et al. described a pre-ignition (PIR)
reaction caused by the fluorination of the alumina shell (Dustin T. Osborne and Pantoya
2007). They describe an interaction of fluorine with the hydroxyl ions boned to the
alumina surface and theorized that the PIR acted to accelerate the reaction by assisting in
the decomposition of Teflon. Iodine has also been shown to react with hydroxyl ions on
the surface of alumina (Flockhart, Liew, and Pink 1974), the surface reaction itself may
contribute to the formation of species that may accelerate the reaction such as I- and IO- ,
F- (Flockhart, Liew, and Pink 1974; Pagni, Kabalka, and Boothe 1988). In 2010, Farley et
al. displayed a bonding of iodine with alumina in a thermo-equilibrium setting between
400 and 500oC by reacting Al2O3 with I2O5 (Farley and Pantoya 2010). In the case of
reaction involving halogenated oxides, the existence of free, unbonded halogen gas could
lead to the halogen bonding to alumina and accelerating the reaction similar to previous
observations (Farley and Pantoya 2010; Saikia et al. 2009; Watson, Pantoya, and Levitas
2008; Gang et al. 2004). The effect may contribute to reduced activation energy and
increased reaction kinetics as seen in the faster burning mixtures. In order to determine
the likelihood of the existence of unbonded halogen gases, a comparative analysis
involving the heat of formation of halides was performed.
The I2O5 adsorption on Al2O3 is strong with bridge bonds forming between
oxygen from I2O5 and surface Al leading to two I-O-Al bonds. It is possible that the
initial composite may not have the ability to overcome the kinetic barrier to form these
bonds. However, the adsorption energy (Eads) is highly exothermic and calculated to be 422.99 kJ/mol for a single I2O5 molecule, or nearly -211.495 kJ/mol per I-O-Al bond
formed. This estimate is important as related I2O5 decomposition fragments such as IO3
and IO2 will also form single I-O-Al bridge bonds for adsorption. In order to compare the
adsorption energy and provide context to this result, adsorption of a water molecule on
iv
Al2O3 is exothermic by -140 kJ/mol on -Al2O3. It is noted that most oxide surfaces are
covered in water in atmospheric conditions. However, I2O5 has a much stronger
adsorption enthalpy and therefore can replace water provided it has enough thermal
energy. At around 200oC this process may start at the interface and proceeds to replace
water. The reactions can be due to I2O5 or HIO3 (hydrated intermediate) and will provide
a similar exothermic response in Al/I2O3 .
On transition from solid-phase reactions to gas-phase reactions, different gasphase species can be present. If we consider that the I2O5 and Al channels are decoupled
except some chemisorption of I2O5 at the Al2O3/I2O5 boundaries, the majority of the
exothermic yield will occur upon thermal decomposition of I2O5; consider the following
steps: I2O5 → IO2 + IO3 → I2O3 + O2. The O2 thus produced can subsequently react with
Al. As Al+O2→Al2O3 reaction happens at much faster rate, the I2O5 can act as accelerant.
The first important result for Al reacting with iodine oxide species as evident
from Table 5 is that most of these reactions are highly exothermic. Secondly, Al can act
as an effective intermediate to aid decomposition of iodine oxide species and move the
oxidation reaction forward. At the same time, steps in which Al-I species are produced
are less exothermic than Al-O species formation. Therefore, Al acts as an intermediate
that aids the production of I and I2 in the gas-phase reaction.
The trends in reactions from Table 4 and Table 5 indicate that the Al/I2O5 has the
following behaviors:
(1) I2O5 provides multiple exothermic channels that contribute to the next
exothermicity;
(2) many of the steps of thermal decomposition produce O2 that sustains the
combustion process and increases reaction temperature; and,
(3) Al-O species are highly stable and a method is shown in which Al-IO or Al-O
species converts to Al2O3 and produces more I2 as the gas-phase product.
These chemical kinetic steps may be the controlling factors accelerating the AlI2O5 and similarly the Al-AgIO3 reactions but cannot be applied to the Al-Ca(IO3)2 and
iv
Al-KClO4 reactions because their onset temperatures are in excess of 200 degrees greater
than for the Al-I2O5 and Al-AgIO3 oxidizers.
iv
CHAPTER IV: THE CONTRIBUTION OF ATMOSPHERIC
OXYGEN TO ENERGY PROPAGATION IN REACTIVE
MATERIALS
.
Introduction
Reactive materials are a class of energetic materials that combine Al fuel with a
solid oxidizer such that the powder mixture is a composite energetic material. When the
particle size of the fuel and oxidizer include one geometric dimension on the nanoscale,
the composite may be referred to as a nano-composite reactive material. Various studies
have examined the effect of atmosphere on aluminum (Al) combustion either as an
aerosolized powder or single particles(Huang et al. 2009; Gill, Badiola, and Dreizin
2010; Beckstead, Liang, and Pudduppakkam 2005; Friedman and Maček 1963).
However, there have been minimal studies examining the effect of environment on flame
speeds of reactive materials. Yet, in practice, reactive materials are often used under
atmospheric conditions such that the role of oxygen in the environment influencing the
combustion should be well understood.
Weismiller and Malchi studied the effects of atmospheric pressure on the flame
speeds of nanometric Al and copper oxide (CuO) mixtures(M. Weismiller et al. 2009).
Their study observed the effect of adding argon, nitrogen or helium up to 15MPa and
noted a shift from convective to conductive heat transfer modes at high pressures(M.
Weismiller et al. 2009). By adding inert gas to the environment, the partial pressure of
oxygen was reduced, thereby creating a low oxygen environment at an elevated pressure.
They showed that in the low oxygen high pressure environment flame speeds decreases
by two orders of magnitude. However, the change in flame speed for this study was
attributed to the pressure change, not environmental gas. In another study, Asay et al.
observed flame speeds of nanometric Al with molybdenum trioxide (MoO3) under
vacuum conditions in a tube(Asay et al. 2004). A vacuum environment is, by definition, a
low oxygen atmosphere, but at a low pressure, and flame speed was found to be
unaffected. Results from these two studies pose questions regarding the role of oxygen in
the environment on combustion at atmospheric conditions.
iv
Two widely studied reactions are Al+CuO and Al with iron (III) oxide (Fe2O3).
Studies have shown that Al+CuO is a gas generating, convectively driven, relatively fast
burning mixture ( i.e., flame speeds greater than 100m/s)(Umbrajkar, Schoenitz, and
Dreizin 2006; Sanders et al. 2007; Bockmon et al. 2005; Jian et al. 2013; Prakash,
McCormick, and Zachariah 2005; M. R. Weismiller, Lee, and Yetter 2011).
Theoretically, the gas production from this mixture could reach 0.54mol per 100g of
powder burned(Fischer and Grubelich 1998). In contrast, Al+Fe2O3 is a conductively
driven reaction with little gas production, e.g., 0.14mol per 100g of powder
burned(Fischer and Grubelich 1998). Also, due to the conductive nature of the reaction,
flame speeds are considerably slower for this formulation than the Al+CuO (i.e., on the
order of 1m/s)(Crane, Pantoya, and Dunn 2010).
Jian et al.(Jian et al. 2013) conducted a study of various powders and correlated
ignition temperatures to oxidizer decomposition temperatures. According to their study,
the mixtures used here ignite when their oxidizer begins to decompose. For example,
CuO decomposes to Cu2O at 1000 K and Fe2O3 decomposes to Fe3O4 at 1400K(Jian et al.
2013). When CuO decomposes, 10% of its mass converts to free oxygen, O, while only
6% of the Fe2O3 results in free oxygen. Therefore, the decomposition of CuO results in
more oxygen available to react with Al, and Umbrajkar et al, showed that oxygen, when
freed from a decomposing oxide, prefers reacting with Al instead of forming
O2(Umbrajkar, Schoenitz, and Dreizin 2006). Also, ignition of nanometric Al in a
gaseous oxygen environment occurs at 1000K.(Huang et al. 2009)
In addition to these two well-studied mixtures, Al+I2O5 and Al+Ca(IO3)2 are also
included in this study. The Al+I2O5 reaction has recently sparked interest due to its highly
convective nature (0.62 mol per 100g burned)(Fischer and Grubelich 1998) and extreme
flame speeds of up to 2000m/s in a confined burn(Martirosyan, Wang, and Luss 2009).
This mixture is rather unique in that the oxidizer interacts with the alumina shell causing
the shell to become an active participant in the reaction instead of just being a diffusion
barrier or heat sink(Farley and Pantoya 2010). Also since part of the combustion products
are I2 gas, this mixture has also been studied as a biocidal agent(B. R. Clark and Pantoya
2010). The Al+Ca(IO3)2 reaction has been found to initiate upon oxidizer decomposition
iv
at 825K and 1025K and Ca(IO3)2 has been found to decompose as follows in Eqs. (38)
and (39). At about 825K, Ca(IO3)2 initially decomposes(K. Stern 1974).
5𝐶𝑎(𝐼𝑂3 )2 → 𝐶𝑎5 (𝐼𝑂6 )2 + 8𝐼 + 18𝑂
iv
(38)
Then at 1050K the decomposition undergoes a second step.
𝐶𝑎5 (𝐼𝑂6 )2 → 5𝐶𝑎𝑂 + 2𝐼 + 7𝑂
(39)
The objective of this study is to examine the influence of atmospheric oxygen on
flame speeds for two convectively dominant reactions: Al+CuO and Al+I2O5, and two
conductively dominant reactions: Al+Ca(IO3)2 and Al+Fe2O3. All reactions initiate when
oxygen is released from the oxide and is available to react with the Al, however each
oxide dissociates at different temperatures either before or after Al melting and when Al
becomes reactive with atmospheric oxygen at 1000K(Huang et al. 2009). The objective
will be accomplished through high speed imaging of reaction propagation defined as
flame speed in controlled oxygen concentration environments.
Methods
Passivated aluminum (Al) with an average diameter of 80nm and an active Al
content of approximately 75wt% was supplied by Novacentrix (formerly
Nanotechnologies), Austin, TX. These particles were used for all tests and mixed with
either CuO or Fe2O3 supplied by Sigma Aldrich, St Louis MO or Alfa Aesar, Ward Hill,
MA, respectively. All particles except I2O5 have a spherical morphology. The average
particle diameter of CuO is 50nm; Ca(IO3)2 is 15 microns; Fe2O3 is 40 microns; and, I2O5
has a flake morphology with the largest dimension no greater than 45 microns. All
reactions were balanced based on oxygen transfer assuming 100% of the oxygen
transferred to the Al.
Individual reactant powders were measured for equivalence ratios of 1 and 1.2,
corresponding to stoichiometric and slightly fuel rich. The powders were then suspended
in 60 cc of hexane and sonicated with a Misonix model S3000 for 70 s with 10 s rest
intervals in order to prevent heating the sample. The solutions were then placed in a glass
tray under a fume hood to allow the hexane to evaporate before the powders were
reclaimed for further experimentation.
Loose powder was loaded into a 0.5 by 13cm notch and distributed evenly with
no packing. By considering the volume of the notch and the mass of the loaded powder,
the bulk density of the sample ranged from 4 to 8% of the theoretical maximum density
(TMD) of the solid mixture meaning 92 to 96% of the volume in the notch is occupied by
iv
air. The mass of the powder varied between 335-720 mg depending on the composite
used to maintain a constant bulk density. For each test, three experiments were performed
to establish repeatability which was found to be the largest source of uncertainty in the
data.
The experimental setup shown in Figure 14 Error! Reference source not
found.consists of a sealed combustion chamber with a blow-off valve to prevent
pressurization. Two roughing pumps are connected to the sealed combustion chamber
and one third of the atmosphere was vacuumed out. Either argon (Ar) or oxygen (O2) was
pumped into the chamber depending on the experiment. This atmospheric cycle was
repeated 4 times for the argon and 6 times for the oxygen as seen in Eqs. (40) and (41).
1 4
21% 𝑂2 (1 − ) = 4% 𝑂2
3
(40)
1 6
(79% 𝑛𝑜𝑛 𝑂2 ) (1 − ) = 7% 𝑛𝑜𝑛 𝑂2 = 93% 𝑂2
3
(41)
Nickel-chromium wire, Ni80-020, from Omega Engineering, Inc was placed at
one end of the notch and heated at 7V by a Variac transformer to provide ignition for the
reaction. Flame speeds were recorded via a Phantom 7 high speed camera at frame rates
ranging from 900 to 90000 frames per second aligned perpendicular to the direction of
flame propagation (Figure 14). National Instruments Vision Builder AI 2010 software
was used to locate the flame front in the high speed images. Initially, a threshold filter
was applied to each frame and the pixel location corresponding to the edge of bright light
was recorded. The pixels were converted to millimeters via calibration and combined
with the frame rate to calculate flame speed.
iv
Figure 14. Top view schematic for open tray flame speed experiments.
Results
Gas generation calculations for each reaction were computed using REAL code
software. The calculations were performed for atmospheric pressure conditions and the
gas production did not vary significantly with stoichiometry such that Table 6 reports the
stoichiometric simulations for comparison.
Table 6. Gas production from REAL Code analyses
Al+CuO Al+I2O5 Al+Ca(IO3)2 Al+Fe2O3
24%
75%
71%
11%
A representative sequence of still frame images captured from the high speed
recording for each mixture is shown in Figure 15. The position versus time data are linear
for each speed measured indicating steady propagation.
iv
Figure 15. A) Al+Ca(IO3)2 at 6.9ms intervals. B) Al+Fe2O3 at 6.9ms intervals. C)
Al+I2O5 at 44µs intervals. D) Al+CuO at 121µs intervals.
Table 7 displays the flame speeds for each reaction under the various conditions.
Flame speeds are organized by equivalence ratio (ϕ) and atmosphere for the four separate
testing conditions. In order to compare flame speeds for the variety of conditions, percent
difference was used as defined in Eq. (42) where FSH is the flame speed under high O2
conditions and FSL is the flame speed for low O2 conditions.
% 𝐼𝑛𝑐𝑟𝑒𝑎𝑠𝑒 =
(42)
|𝐹𝑆𝐻 − 𝐹𝑆𝐿|
∗ 100%
𝐹𝑆𝐿
Table 7. Summary of flame speed results in units of m/s.
Al+I2O5
High O2
Low O2
%
Increase
Al+CuO
Al+Ca(IO3)2
ϕ=1.0
ϕ=1.2
ϕ=1.0
ϕ=1.2
ϕ=1.0
840±4
700±3
224±
408±1
2.9±0.0
0
5
4
5
4
770±2
670±6
204±
320±2
0
0
9
9%
4%
10%
Al+Fe2O3
ϕ=1.2
ϕ=1.0
4.5±0.7
3.4±0.1
2.2±0.0
2.9±0.0
1.7±0.0
1.8±0.
5
5
1
4
1
30%
30%
55%
100%
200%
iv
ϕ=1.2
5.4±0.
2
Discussion
The flame speeds shown in Table 7 suggests combustion of solid fuel-oxidizer
mixtures in an oxygen rich environment occur in different ways. Based on the
literature(Umbrajkar, Schoenitz, and Dreizin 2006; J. H. Stern and Passchier 1962; Jian et
al. 2013; Fischer and Grubelich 1998) and observations from Table 7, Figure 16
illustrates a flow chart describing the mechanisms promoting flame speed. For solid
oxidizers with decomposition temperatures below 1000K, the first stage is the
decomposition of the metal oxide. Next, stage two is marked by the ignition of Al with
oxygen freed from the solid oxidizer. The heat from this reaction will increase
temperatures and provide energy to dissociate atmospheric O2 to O. As Al particles reach
1000K, stage three is designated by reactions of Al with atmospheric oxygen. The heat
generated from these reactions increases local temperatures to over 3000K(Dreizin et al.
2011), such that stage four occurs when the rest of the metal oxide decomposes and
provides additional energy to dissociate atmospheric O2 to O. it is noted that the heat of
formation for O from O2 is endothermic and 249 kJ/mol, while the heat of formation of
Al2O3 is -1670 kJ/mol(Cox, Wagman, and Medvedev 1987). In stage five, the remaining
Al then reacts with the available oxygen from the solid decomposition and the
atmosphere. Finally, with sufficient oxygen available, reduced metal products will then
react with remaining atmospheric oxygen. All of the stages involved in producing the
convectively dominant fastest flame speeds are shown graphically in Figure 16.
In a low oxygen environment, the oxide will still decompose such that oxygen is
available to initiate combustion by 1000K. Stages three-five continue to occur, with less
atmospheric oxygen participating, less convective influence and slightly reduced flame
speeds. A slower conductively driven reaction such as Al+Fe2O3 follows the flow chart as
well, but with higher decomposition temperatures, the mechanism limits stages one and
two.
iv
Stage 1: Decomposition of
solid oxidizer below 1000K
Stage 2: Reaction of Al with
dissociated oxygen
Stage 3: Reaction of Al with
atmospheric O2 at 1000K
Stage 4: Further
decomposition/reaction
>1000K
Stage 5: Reaction of
reduced metal products
Figure 16. Flowchart of fast convective reactions
Iodine pentoxide
Iodine pentoxide decomposes at 663K into IO3 and IO2(Farley and Pantoya 2010).
These two compounds bond exothermically to the passivating Al2O3 shell on the nAl
particles. The IO3 and IO2 compounds then further disassociate at 730K into I and O gas
reacting with Al to form Al2O3 and I2. At this point, a fully decomposed I2O5 molecule
produces 100% total gas of which 24% is oxygen by mass. From this sequence of stages
in Figure 16, nearly all of the oxygen from I2O5 becomes available to react in stage one.
As a result of the low decomposition temperature and complete gas disassociation of the
I2O5 particle, this reaction is highly convective with large pressure increases and
extremely high flame speeds (670-840 m/s) regardless of atmospheric oxygen
concentration. Adding or removing atmospheric oxygen into this reaction affect heat
generation by reacting with Al that is left unreacted. As a result, changing the
atmospheric oxygen content from 4-93% increases the flame speed by only 9% for the
stoichiometric mixture. It should be noted that the flame speeds for the fuel rich (ϕ=1.2)
mixtures have slower flame speeds than the stoichiometric mixtures. This is possibly due
iv
to a reduced amount of IO2 and IO3 radicals to bond to the alumina layer reducing the
early onset of heat generation (stages one and two) and reducing the flame speed.
Copper oxide
In view of the five stage approach to understanding atmospheric oxygen effects
and flame speeds (Figure 16), while Al+I2O5 is largely a stage one and two reaction,
Al+CuO utilizes all five stages with an emphasis on stages one-three. As previously
discussed, the initial decomposition of CuO→Cu2O+O begins at 1000K, releasing 50%
of the oxygen. This temperature also corresponds to Al particles initially reacting with
atmospheric O2 and published ignition values for Al+CuO(Jian et al. 2013). In other
words, Al+CuO in a sense, compresses stages one-three into a narrow temperature band
at 1000K. Afterwards, Cu2O decomposes and reacts in stage four and Cu reacts with
atmospheric O2 in stage five to help propagate the reaction. This reaction has long been
considered convective producing a lot of gas (25% from thermal equilibrium code) with
very high flame speeds (200-400 m/s). Since a large portion of this reaction begins at
lower temperatures (stage three), the reaction is less affected by atmospheric oxygen (1030%) than Al+Fe2O3 and Al+Ca(IO3)2, but more affected than Al+I2O5. Increasing the
content of oxygen surrounding the mixture increases the effects of stages three-five and
provide more heat and product gas to convect the reaction forward. Higher equivalence
ratios (ϕ=1.2) result in slightly higher gas production and higher flame speeds. Increasing
oxygen concentration for ϕ=1.2 mixtures results in the largest increase in flame speed for
this mixture due to the exaggerated effect of the initial stage since more Al is available to
react to the O2.
Calcium Iodate
As a result of the two stage decomposition as seen in Eqs. (38) and (39),
this reaction follows the stages in a similar fashion to Al+CuO. Roughly 50% of the
oxygen content is released right before the Al powder reacts at 1000K. Then, as the
temperature increases, stage four occurs with the second stage of the decomposition. If
this five stage process explaining the effect of atmospheric oxygen was all encompassing,
then the flame speed for Al+CuO and Al+Ca(IO3)2 would be expected to be the same.
However, Ca(IO3)2 and Ca5(IO6)2 (Hf = 2.57 and 5.92 kJ/g)(K. Stern 1974) are much
iv
more stable compounds than CuO and Cu2O (Hf = 1.95 and 1.17 kJ/g)(Sottery 1985). As
a result, even though the decompositions occur at similar temperatures and follow similar
steps, the heat absorbed in order to decompose both materials are drastically different. As
a result flame speeds are reduced by the decrease in latent heating to the surrounding
material. By increasing atmospheric oxygen, more heat can be generated from reacting Al
to help compensate for the large amount of heat absorbed in the decomposition process.
As a result, flame speeds are more sharply increased as compared to Al+CuO (30-55%).
Combining more Al (ϕ=1.2) and a high level of O2 results in the highest heat output and
similarly highest flame speed.
Iron oxide
The decomposition of Fe2O3 occurs at 1400K such that the initial phase for
Al+Fe2O3 favors Al reacting with atmospheric oxygen to produce energy promoting
Fe2O3 decomposition. As a result, this reaction is largely a stage three-five reaction. The
role of atmospheric oxygen is more critical with this mixture in this first phase of reaction
because the solid oxidizer provides less oxygen for Al reaction. The difference in initial
oxidizer decomposition between Al+CuO and Al+ Fe2O3 reactions may be key to
understanding the dramatic increase in flame speed for Al+ Fe2O3. While Al+CuO has a
large concentration of oxygen available at stage three, Al+Fe2O3 only has atmospheric
oxygen available under half of the stages. When atmospheric oxygen is not present, the
reaction does not initiate until stage four. However, with the presence of atmospheric
oxygen, Al can react in stage three to accelerate decomposition of Fe2O3. None of the
other reactions are as largely affected by atmospheric oxygen and as a result, Al+Fe2O3 is
impacted the most with flame speed increases of 100%. Adding extra Al increases the
initial reactivity under all conditions, but combined with a high O2 atmosphere, the
increased emphasis of the stage three reaction increases flame speeds by 200%.
Assuming that the relationship between the temperatures of oxidizer
decomposition and reaction with atmospheric oxygen are key to the sensitivity of a
reaction to atmospheric oxygen, a non-dimensional sensitivity factor (S) is introduced.
The sensitivity factor should also include information on reaction pressures that can
affect the ability of atmospheric oxygen to come in contact with the Al particles and
iv
ignite. Generally, fast reactions (flame speed (FS) >100m/s) are more convective with a
greater pressure differential created by the high speed. In this way the sensitivity factor
also includes convective effects by comparing the log10(FS) to the log10(100).
𝑆=
𝑇𝑑 − 1000 log10 (100) − log10 (FS)
+
1000
log10 (100)
(43)
In Eq (43), Td is oxidizer decomposition temperature in Kelvin. The sensitivity
factor includes a combination of the thermal decomposition sensitivity initiating the
reaction and the convective sensitivity for the ability of atmospheric oxygen to contribute
to the reaction once ignition has occurred. The revised sensitivity factors are shown in
Figure 17 alongside the data symbol for the reaction. Figure 17 provides a relationship
for the sensitivity of the reaction to influences of atmospheric oxygen that is consistent
with the diverse mixtures studied here. Future work can further affirm this predictor and
assess the role of atmospheric oxygen on flame speed.
120
Al+I2O5
Al+CuO
Increase in FS (%)
100
1.2, 100
Al+Ca(IO3)2
80
Al+Fe2O3
60
40
0.62, 30
20
-0.72, 9
-0.16, 10
0
-1
-0.5
0
0.5
1
1.5
Sensitivity Factor
Figure 17. Percent increase in flame speed as a function of sensitivity factor
iv
CONCLUSIONS
Differential scanning calorimetry (DSC) and thermogravimetric (TG)
analysis of I2O5, I2O5/Al and I2O5/Al2O3 mixtures in argon show a scale dependent
reaction based on the size of the Al powder. While larger scale micron powders show
little to no reactivity, nanometer scale passivated Al powders undergo a complex
multistep reaction when mixed with I2O5. The reaction commences upon I2O5
decomposition and is triggered by I- adsorption into the Al2O3 passivation shell, further
reactions between the iodine and oxygen gas and Al in the solid phase is then observed.
This study found no correlation between heat of combustion or mass loss and
flame speed in open tray burns. Instead, apparent activation energy correlates with flame
speeds by orders of magnitude suggesting flame speeds are attributed to diffusion limited
but kinetically controlled reactions. In order to understand why the activation energy was
lower for these reactions, studies of the free halogen gas were conducted showing
intermediate chemistries contribute toward heightened levels of exothermicity. Surface
chemistry associated with iodine interaction with the alumina shell is considered a first
step in the reaction mechanism. Intermediate chemistries between iodine species and
aluminum are the controlling accelerant mechanism for Al-I2O5 and Al-AgIO3 reactions.
In contrast iodine radicals are not as easily released in Al-Ca(IO3)2 and Al-KClO4 with
onset temperatures in excess of 200 degrees higher for these mixtures.
Theoretical analyses show that almost all of the intermediate reactions between Al
and iodine species are highly exothermic. Also, Al can act as an effective intermediate to
aid decomposition of iodine oxide species and aid in the production of I and I2 in the gas
phase. In summary the chemical kinetics responsible for the low activation energies and
high flame speeds observed for Al-I2O5 and Al-AgIO3 are: (1) I2O5 and IO3 provide
multiple exothermic channels that contribute to the next exothermicity; (2) many of the
oxidizer thermal decomposition steps produce O2 that sustains the combustion process
and increases reaction temperature; and, (3) Al-O species are highly stable and a method
in which Al-IO or Al-O species converts to Al2O3 and produces more I2 as the gas-phase
product has been presented. Results for these experimental and theoretical investigations
elucidate the factors controlling reactivity of nano-Al energetic composites containing
iv
halogenated oxide species. Theoretical chemical influences describe in terms of reaction
kinetics explain the observed apparent activation energy and flame speed measurements.
Flame speeds were measured for four composite reactions. All reactions are
highly gas generating except Al+Fe2O3. However, Al+Ca(IO3)2 has a flame speed of only
2.5 m/s implying a primarily conductively dominant reaction despite high gas generation.
Results show increasing atmospheric oxygen has a much more dramatic effect on the
flame speed of conductive reactions, Al+Fe2O3 with a 100-200% increase and
Al+Ca(IO3)2 with a 30-55% increase. For the convective reactions, Al+CuO increased by
only 10-30%, and Al+I2O5 by 4-9%. This behavior is the basis for a model that describes
a sensitivity factor based on solid oxidizer decomposition temperature and flame speed
that accounts for oxygen availability and thus how the flame speed will be affected by the
presence of atmospheric oxygen. Earlier onset of solid oxidizer decomposition
temperatures and higher flame speeds cause reactions to be less sensitive to atmospheric
oxygen concentration. In general, these results suggest atmospheric oxygen plays a
greater role enhancing flame speeds for conductively dominant reactions. Since the boost
to flame speed is limited by Al availability, adding more fuel increases flame speed and
results show that increasing equivalence ratio from 1.0 to 1.2 produces accelerated flame
speeds.
iv
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