Download Kinetics and Equilibrium of the Reversible Formic Acid

P R E P R I N T – ICPWS XV
Berlin, September 8–11, 2008
Kinetics and Equilibrium of the Reversible Formic Acid Decomposition in Hot Water
Yoshiro Yasaka, Ken Yoshida, Chihiro Wakai, Nobuyuki Matubayasi, and Masaru Nakahara
Institute for Chemical Research, Kyoto University
Email :[email protected]
During the past decades, a variety of technology has been proposed for the application of
hydrogen energy. The fundamental problems are its storage and the transportation, and they
have not yet been solved. This article explains how to use formic acid as a fuel; the advantages
of carbonaceous fuels and hydrogen are fulfilled by formic acid. The strategy is shown here how
to produce and utilize the formic-acid fuel on the basis of the detailed kinetic and equilibrium
studies on the reversible formic acid decomposition in hot water,
Introduction
To overcome the environmental issues caused
by extensive and rapid fuel consumption and
hazardous chemical processes in industry is one of
the greatest subjects for the chemistry of the 21st
century. Hot water is a promising candidate that
would play a key role in recovering the sustainable
energy and material (carbon) cycles on earth. More
and more attention is paid for the hydrogen cycle;
see Figure 1. The most attractive potential of hot
water for this application is that hot water can
dissolve various chemical species including C1
(carbon one) molecules, plastics and biomass, and
can induce unique hidden reactions among them in
the absence of metal catalysts, thus allowing us to
convert less useful materials into more useful ones
[1].
The atmospheric carbon dioxide issue, which is
now the most serious and urgent environmental
problem all over the world, has its root in human
activities interfering with the natural carbon cycle.
The accumulation of carbon dioxide in the
atmosphere can be interpreted by considering a
large gap between the oxidation rate of fossil fuels
by human beings (fast step) and the reduction rate
of carbon dioxide by nature (slow step); see Figure
2. We are going to completely exhaust the millionsof-years-old fossil fuels in just a few centuries. The
carbon dioxide problem is strongly related with the
energy supply problem because the oxidation
process is accompanied by useful energy of up to
800 kJ per one mole of carbon (44 g CO2).
Fortunately, we receive at any moment plenty of
solar energy in comparison to the energy we need
for our daily activities. The solar energy we receive
on earth amounts to 5730 TW, while the global
energy consumption by human beings amounts to
the one-thousandth [2]. A key point of the energy
problem is the fact that so far we make use of only
a small portion of solar energy. We have just started
a large-scale fixation of solar energy in the form of
a controversial man-made fuel called bioethanol.
Can we take advantage of the unique potential of
Figure 1. Carbon and hydrogen cycles for energy through
formic acid.
the gaseous species. This is a kind of chemical
evolution.
The carbonaceous energetic compounds
produced from solar energy are very convenient to
store because they are in the liquid or solid state.
Typical examples are such fossil fuels as oil and
coal. When they are burned, however, they are not
so clean as hydrogen energy because of the
accompanying production of nitrogen oxides,
sulphur oxides, and carbon micro particles in
exhaust fumes. Hydrogen is ultimately a clean
energy resource but is extremely difficult to store
because of the explosivity and lowest condensation
temperature as seen in Figure 3. This is a huge
obstacle we are faced with in developing hydrogen
energy technology. The WGS reaction via formic
acid is a key reaction that allows us to take
advantage of the benefits of both types of fuels; we
can store and transport formic acid just as we store
petroleum and can use it as if it were hydrogen by
the application of the second step of reaction (2).
Formic acid is neither explosive, flammable nor
very volatile; as a drawback it is weakly acidic. For
these reasons, formic acid can be a potential manmade fuel alternative to bioethanol.
Any weakly reducing carbonaceous materials
can be used for the input energy for the generation
of this new fuel from water. When we succeed, we
need not rely upon corn or sugarcane for bioethanol.
The strategy for the synthesis is shown in Figure 4.
As can be seen, the partial combustion of waste
carbonaceous materials yields carbon monoxide,
hot water to go beyond bioethanol? We must reduce
CO2 emission by 60-80 % by the year of 2050 (EU
summit meeting on May 2007).
Formic Acid as a Clean Fuel
The water-gas-shift (WGS) reaction [3]
CO + H2O
CO2 + H2
(1)
is one of the oldest hydrothermal reactions studied.
It was industrially developed in the early 20th
century for the steam-reforming process. However,
the reaction mechanism under non-catalytic
conditions has not been established for a long time.
The first attempt to detect formic acid as an
intermediate of the WGS reaction was done by Rice
et al. using a flow system made of metal in
supercritical water [4]. After their unsuccessful
results, we have carried out series of studies in
much milder reaction conditions [5], and after all
succeeded in detecting formic acid produced from
carbon monoxide in sub-critical water, which then
decomposes into carbon dioxide and hydrogen [5c].
The WGS reaction is then rediscovered as a
reversible transformation process between carboncontained chemical energy (HCOOH) and
hydrogen-contained chemical energy (H2):
CO + H2O
HCOOH
CO2 + H2 (2)
The equilibrium can be shifted to either side
(CO+H2O or
H2+CO2)
by
tuning
the
thermodynamic conditions. Furthermore, we have
shown that the reaction can be kinetically
controlled in order to synthesize formic acid from
CO2
Sun
breathing
Artificial
Fuels
in Future
photosynthesis
fossilization
Life
Fossile Fuels
Figure 2. Materials and energy flow on earth.
2
(a)
cell technology (electricity
H2 ) and the
reversible reaction (2) make this strategy very
hopeful.
The grand design of the formic acid fuel drafted
above requires fine control of the WGS reaction via
formic acid. The following sections introduce the
basic hydrothermal chemistry of formic acid [5].
CO + H2O
HCOOH
Formic Acid
Fuel
CO2 + H2
Formic Acid Decomposition in Relation to the
WGS Reaction
100 °C
(b)
H2
N2
In recent years, [5-10] formic acid is found to
decompose without catalysts in the following ways:
HCOOH
CO +H2O
(decarbonylation)
(3)
HCOOH
CO2 +H2
(decarboxylation)
(4)
Using NMR, we have found that the
decarbonylation is reversible and that carbon
monoxide can be directly converted into formic
acid in hot water [5c]. We also made a theoretical
analysis on the Gibbs energies of the species
(HCOOH, CO, CO2, H2, and H2O) involved in
reactions (3) and (4) [5d]. The reversibility and the
coupling of the reactions given by eqs 3 and 4
clearly indicate that formic acid exists as an
intermediate in the WGS reaction as in eq 2.
To comprehensively understand and control
the WGS reaction, it is essential to determine the
rate and equilibrium constants for the elemental
steps shown in eqs 3 and 4. The kinetics and
equilibrium study was hampered by the following:
(i) the reaction order of the reactions (3) and (4) has
not yet been established, (ii) these reactions can be
catalyzed by metals used for reaction cells [5b,7,10],
(iii) although the non-catalytic decarbonylation can
be dominant in hot water, the reaction rate is very
slow (the time scale of days to weeks below ~200
°C), and (iv) it is rather difficult to quantitatively
analyze such main gaseous products as CO, CO2,
and H2 in the heterogeneous system because of their
distribution between the gas and liquid phases. Here
we show that the most of these difficulties can be
overcome by studying the decomposition of the
formic acid in the presence of acid but in the
absence of metal surface, because under such
condions the two competitive decomposition
pathways in eqs. 3 and 4 can be separated to obtain
each reaction rate constant. The difficulty in the
quantification of the products can be overcome by
applying 1H and 13C NMR spectroscopy to the
liquid and gas phases of the sealed reactor, with the
carbon mass balance confirmed.
101 °C
CH4
H2O
HCOOH
-160 °C
-196 °C
-253 °C
Figure 3. The comparison of the condensation
temperature of fuels and related compounds (b)
involved in the formic-acid formation reaction
(a).
which is then hydrothermally converted into formic
acid. The strategy as a whole is equivalent to the
fixation of the solar energy, since the biomass used
is grown by sunlight. The growth of plantlife is the
most efficient and cheapest way as the first step of
the solar energy. At the present stage, in fact, the
global energy fixation rate by plants is estimated to
be ~10 times faster than the global energy
consumption rate by human activities [2].
Another energy source for the production of
formic acid fuel is the electricity from power plant
working with natural energy. Although wind power
and solar power plants show a high performance in
extracting natural energy, the output is unstable due
to the temporal fluctuation and discontinuity of the
given natural energy. This is a drawback and
significantly reduces the overall efficiency of the
power plants. It is thus essential to develop a way
for the large-scale temporary energy storage to
balance the energy supply and demand. The use of
a formic-acid chemical tank is one of possibilities
for this energy storage. The combination of the fuel
3
Partial combustion
biomass
waste plastic
H2O
CO
Formic Acid
Fuel
WGS Reaction
Fuel Cell
Natural energy
power plant
H2
CO2
Figure 4. Two ways of producing formic acid fuel.
that the main product was CO2 [6]. Other authors
also reported at higher concentrations (0.13 M and
1 M) that the decarboxylation is the main
decomposition path [7,8]. McCollom et al.
examined the decarboxylation of formic acid using
a gold bag reactor with and without minerals and
clarified that the catalytic effect is not given by
minerals but by the transition metals except for gold
[10]. They also observed the formation of formate
ion from CO2 and H2 in basic conditions, which
indicates the reversibility of the decarboxylation.
In the following part, we review our most
recent studies [5] and discuss the results in relation
to the formic acid fuel technology.
Hydrothermal chemistry receives further
attention because a number of organic reactions in
hot water involve formic acid as a key product or
intermediate in hot water [5,11]. We need to
establish the kinetics and mechanisms of the
hydrothermal formation and decomposition
reactions of formic acid, in view of the following
three processes [11]. (1) Formic acid, as an
aldehyde (C1, single-carbon contained), is involved
in the non-catalytic cross-disproportionations of
aldehydes; an aldehyde is reduced to the
corresponding alcohol and formic acid is oxidized
to carbonic acid or carbon dioxide. (2) Formic acid
reacts with formaldehyde (C1 aldehyde) in the
presence of acid to form a C2 compound, glycolic
acid, through C-C bond formation as a chemical
evolution process. (3) Carbon monoxide, which is
in a hydrothermal equilibrium with formic acid, is
released through the decarbonylation of aldehydes
that can be generated by the heterolytic
fragmentation of ethers.
The decomposition of formic acid has been
studied for a long time, but much attention was paid
to the catalytic effect of the metal surface. Our
group and Bröll et al. reported the dominance of the
decarbonylation path in the decomposition at low
temperatures without kinetic analysis [5a,5b,9]. The
decarboxylation pathway was much more
quantitatively studied [6-10], but the reaction rate
was obtained mainly under the surface catalytic
effect of metals. Savage and co-workers studied the
decomposition of formic acid in a flow reactor
made of Hastelloy C-276 at 0.02 M (mol dm-3) in
the temperature range of 320-500 °C and showed
Experimental
A quartz tube reactor was used in this
study as shown in Figure 5. The reaction solution
was introduced to occupy a specified portion
(controlled with an accuracy of 1%) of the tube
reactor. The volume (length, lTliq) ratio of the liquid
phase to the total (L) is hereafter denoted as the
filling factor:
ρT =
liq
the volume of the liquid phase lT
=
the total volume of the reactor
L
(5)
where the suffix T represents the temperature. Note
that the filling factor ρT at a reaction temperature T
is larger than that at room temperature 30 °C, ρr,
due to the expansion of water when ρr > 0.32 (at
densities higher than the critical density of water
0.32 g cm-3). The gaseous reaction product is in a
dissolution equilibrium at each temperature. The
4
value of ρT can be calculated from ρr by using the
tabulated density of water along the saturation
curve [12].
The quartz tube reactor containing the reaction
mixture was heated for a specified reaction time in
a programmable electric furnace kept at high
reaction temperature (170-330 °C) with an accuracy
of 1 °C. The reaction proceeds in an equilibrium
with the gases. After the sample was quenched by
air, 1H and 1H-decoupled 13C NMR spectra were
observed at 30 °C with a NMR spectrometer (JEOL
ECA400, magnetic field strength of 9.4 T). Such
gaseous products as CO, CO2, and H2 coexist in the
liquid and gas phases and the NMR spectra were
measured for each phase by the method described
elsewhere [11a]. The time evolution of the
concentrations was measured by repeating the
above procedure.
The concentrations of all species (HCOOH,
CO, CO2, and H2) were quantified on the basis of
their integrated peak intensity. For example, the
concentration of formic acid [HCOOH] was
estimated by the 1H peak intensity relative to that of
solvent H2O (55.5 M) in the liquid phase. Then the
fraction of residual formic acid, denoted as xHCOOH,
is expressed by
x HCOOH =
of the gaseous products such as CO was summed
over the liquid and gas phases. For the computation
we used the following equation:
[CO]liq + [CO] gas ×
xCO =
1 − ρr
ρr
[HCOOH]0
(7)
Here, [CO] and [CO] are the concentrations of
CO in the liquid and gas phases, respectively,
obtained at 30 °C. The same procedure was used for
CO2. No other side-reactions took place because the
carbon mass balance was maintained during the
reaction within 10%. The source of error is mainly
the absence of an internal reference in the estimate
of the gaseous products. The maintained mass
balance also indicates that CO and CO2 do not leak
through the reactor walls. However, H2 leaked
through the quartz walls very slowly as pointed out
in the original paper [5f].
Because the reaction is carried out in
heterogeneous (sub-critical) conditions, the gaseous
reaction products are present in both phases. For
this reason, we need to take into account the
dissolution equilibrium of the gases at the reaction
temperature when determining the equilibrium
constant for the decarbonylation defined as:
liq
[ HCOOH ]
[ HCOOH ]0
K CO =
(6)
where [HCOOH]0 is the initial concentration of
formic acid. The yields of gaseous products were
determined in a similar way but by using an exteral
reference for integrated intensity. The concentration
gas
[CO ] liq
[HCOOH ] liq
(8)
eq
where the symbol |eq represents the value evaluated
at equilibrium. See the following sections in detail.
Figure 5. The experimental procedure used in this study and the accompanying change of the filling
factor at each step. The illustration shows the reactor tube (a) at room temperature before the
initiation of the reaction, (b) during the reaction at a high temperature, and (c) after being rapidly
cooled to room temperature. The filling factor at the reaction temperature is larger than that at room
temperature due to the decrease of water density at elevated temperature for ρr > 0.32. For (b) and
(c), there are also shown the reversible decarbonylation of formic acid and the dissolution
equilibrium of carbon monoxide in the reactor. Note that the decarbonylation reaction is quenched
in (c).
5
Decarbonylation Kinetics
Decarbonylation Equilibrium
Under the catalytic effect of metals, the
decarbonylation pathway can be controlled to be a
minor one. However, we have found that in the
absence of the metal catalysts, formic acid
decarbonylates
much
faster
than
the
decarboxylation. When the HCl concentration is
larger than 2.0×10-2 mol kg-1, the decarbonylation is
exceedingly faster (more than 50 times) than the
decarboxylation and become practically the only
decomposition pathway. In this situation, the initial
rate of carbon-monoxide production is equal to that
of the formic-acid decomposition. We can assume
the rate law of the decarbonylation as
We have determined the equilibrium constant of
the decarbonylation of formic acid in hot water in
the temperature range of 170-280 °C. We have
separated the time scale of the decarbonylation
(fast) from that of the decarboxylation (slow). What
is directly obtainable from the yield analysis is not
KCO but a quantity denoted as the decarbonylation
quotient QCO:
v+1 = −
QCO =
d[ HCOOH ]
b
= k +1[H + ]a [ HCOOH]0
dt
0
log (k+1/ mol-1 kg s-1)
-4.82
-4.22
-3.71
-3.29
-2.98
-2.75
-2.43
-1.91
-1.20
x HCOOH
eq
(11)
where xHCOOH and xCO are provided by eqs 6 and 7,
respectively. As the reaction takes place in
subcritical water along the saturation curve, not all
of the CO produced by the decarbonylation remains
in the liquid phase, but a considerable amount of
CO is accommodated in the gas phase of the sample
at high pressures. Hence QCO is larger than KCO. To
determine KCO we need to calculate the fraction, β,
of CO distributed in the liquid phase.
(9)
Here v+1 is the reaction velocity, k+1 is the rate
constant, [HCOOH]0 is the initial concentration of
formic acid. The parameters a and b are the reaction
orders with respect to the proton (HCl) and formic
acid, respectively. As is shown in ref [5f], they
were determined as unity. Thus we have
v+1=k+1[H+][HCOOH]
(10)
This rate law for the decarbonylation is considered
to be valid also in the absence of acid catalysts or in
basic conditions; the decomposition of sodium
formate produced no detectable carbon monoxide.
On the basis of the reaction rate law (10), we
determined the second-order rate constant k+1 in the
temperature range of 170-280 °C on the liquid
branch of the saturation curve. The results are
summarized in Table 1.
T/C
170
185
200
210
220
230
240
250
260
280
300
x CO
β=
moles of CO in the liquid phase
total moles of CO
(12)
Then the relation of QCO and KCO is written as:
K CO =
βxCO
xHCOOH
= βQCO
(13)
where formic acid is considered to be present only
in the liquid phase.
The value of β is dependent on the solubility
of CO at the reaction temperature and the filling
factor defined by eq 5. The value of β has been
determined on the basis of the following
assumptions: (1) the dissolution of the gases is
sufficiently faster than the decarbonylation and the
decarboxylation. (2) The Henry law holds for the
dissolution of CO into hot water. (3) The equation
of state for the ideal gas holds for CO in the gas
phase.
The Henry’s constant kH is the mole ratio of
liquid H2O to CO dissolved in hot water which is in
a dissolution equilibrium with hypothetical ideal
gas of CO at 1 MPa. When the partial pressure of
CO is given by pCO/MPa, we have the molar
(volume) concentration of dissolved CO as
KCO
0.15
0.33
0.44
0.80
1.2
4.2
Table 1. The rate constant k+1 and the equilibrium
constant KCO for the decarbonylation, and the acid
dissociation constant Ka of formic acid in the
temperature range of 170-330 °C on the liquid
branch of the saturation curve.
[CO]liq =
6
pCO dT
kH
(14)
where dT is the molar concentration of water at the
reaction temperature T on the saturation curve. The
gas-phase molar concentration of CO is expressed
by
[CO]gas =
pCO
RT
Decarboxylation Kinetics
The kinetics of decarboxylation were found to
be more complicated than those of the
decarbonylation. The proton concentration does not
linearly increase the decarobxyration rate. We
measured the decarboxylation rate of formic acid in
the presence of HCl as well as the decarboxylation
rate of sodium formate. The results in Table 2 are
summarized on the following rate law:
(15)
according to the ideal gas law; R is the gas constant.
Using eqs 8 and 11-15, we can evaluate the value of
β as
d[CO 2 ]
= v+ 2 = k [ HCOO − ]
dt
= k acid [ HCOOH ] + k base [ HCOO − ]
[CO ]liq ρ T
β=
[CO ]liq ρ T + [CO ]gas (1 − ρ T )
=
ρT
(19)
K D (CO) (1 − ρ T ) + ρ T
(16)
This rate law was previously postulated by
Maiella et al. [8] and McCollom et al. [10]. The
precision of the rate constants in Table 2 is
relatively low (~10 %) because the surface of the
quartz tube is corroded in basic conditions and the
reaction analysis can be carried out only at an early
stage. Two rate constants kacid and kbase are of the
same order. This indicates that the decarboxylation
rate changes only weakly with pH in contrast to the
decarbonylation. The magnitude relation between
kacid and kbase changes at ~280 °C, and at the higher
temperature of ~ 330 °C, kbase overwhelms kacid.
Taking advantage of the different pH
dependence
of
the
decarbonylation
and
decarboxylation rates, the reaction path control of
the hydrothermal decomposition of formic acid can
be realized. The decarbonylation is the major path
in acidic conditions due to the catalytic effect of the
proton, while the decarboxylation becomes
dominant when pH > 4.
where the liquid-gas distribution constant is
introduced by
K D (CO ) =
kH
[CO ]gas
=
RTd T
[CO ]liq
(17)
and ρT is the filling factor at the reaction
temperature. The Henry’s constant kH taken from
ref [13] and the molar concentration dT of pure
water tabulated in ref [12] are used to evaluate
KD(CO). As the temperature increases along the
saturation curve, liquid water is expanded and CO
becomes more soluble in water in the temperature
range studied here; KD(CO) = 29.6, 13.8, and 8.4 at
170, 240, and 280 °C, respectively. The value of ln
KD(CO) and its temperature dependence are in good
agreement with a previous computational study [5d].
Now the equilibrium constant KCO is explicitly
written in terms of the directly measurable
quantities QCO and ρT as
K CO =
ρT
K D (CO)(1 − ρT ) + ρT
T / °C
200
230
240
260
280
300
330
QCO
(18)
On the basis of this equation, we can determine KCO
by measuring the equilibrium mole fraction
xHCOOH(eq) of formic acid under a given filling factor
ρT. We can evaluate QCO in eq 18 from eq 11;
xCO(eq) = 1-xHCOOH(eq). The equilibrium constant KCO
as the average of two measurements (13
measurements for 240 °C) of different filling
factors are listed in Table 1 at 170-300 °C.
log (kacid / s-1)
-6.6
-5.7
-6.1
-5.9
-5.4
-4.6
-4.1
log (kbase / s-1)
-6.4
-5.5
-4.9
-4.3
-3.4
Table 2. The rate constants kacid and kbase in the
temperature range of 200-330 °C on the liquid
branch of the saturation curve.
7
Decarboxylation Equilibrium
Control of WGS Reaction for Formic Acid Fuel.
Here we discuss the decarboxylation
equilibrium. We followed the decarboxylation for a
sufficiently long time and the time evolution is
shown in Table 3 at 330 °C. As can be seen, the fast
decarbonylation attains equilibrium within an hour,
and this pre-equilibrium is maintained (xCO/xHCOOH
~ 5) during the course of the slow decarboxylation.
The parallel decrease of formic acid and carbon
monoxide almost stops at ~20 h and this indicates
that the decarboxylation equilibrium is attained.
The observed yield of hydrogen is slightly smaller
than the 1 equiv. of CO2, and this probably
indicates the leakage of hydrogen through the
reactor walls (the time scale of the leakage is
estimated to be ~100 h from long-time experiments).
The equilibrium constant KCO2 of the
decarboxylation is defined by
In the previous sections, we have established the
kinetics and equilibrium of the decarbonylation and
the decarboxylation of formic acid which constitute
the hydrothermal water-gas-shift (WGS) reaction.
Based on the above analysis, here we propose a
strategy how we can store energy in the form of
formic acid fuel.
As we have mentioned in the introduction the
energy source for the production of formic acid fuel
is carbon monoxide obtained from any
carbonaceous materials. To convert carbon
monoxide into formic acid, it is necessary to
suppress the WGS reaction at the stage of the
formic acid intermediate; see Figure 6a. The above
kinetic study on the decomposition of formic acid
has proven that in the presence of HCl, the
decarbonylation has a much shorter time scale than
does the decarboxylation. This is an ideal situation
for the formic acid fuel production. The conversion
ratio of CO into formic acid is, in turn, determined
by the equilibrium constant KCO of the
decarbonylation. The expected yield of formic acid
xHCOOH was calculated using the data in Table 1 and
is shown in Figure 7 as a function of temperature
and the filling factor ρT at the reaction temperature.
Better formic acid yield is obtained at a lower
temperature and higher filling factor due to the
favorable KCO and the increased partition of carbon
monoxide into water.
K CO2 =
[CO 2 ]liq [ H 2 ]liq
[HCOOH ]liq
eq
(20)
By taking account of the liquid-gas distribution of
H2 and CO2, the value of KCO2 at 330 °C is
estimated to be 1×102 mol kg-1. Our result shows
that the decarboxylation equilibrium strongly favors
the CO2 side. This is consistent with the
thermodynamic calculations by McCollom et al.
[10], and our statistical-mechanical calculations
[5d].
xHCOOH
1
0.27
0.04
0.014
0.0067
xCO
0
0.49
0.30
0.10
0.056
0.036
0.034
0.039
xCO2
0
0.25
0.66
0.89
0.94
0.96
0.96
0.96
(a)
xH2
0
0.09
0.66
0.79
0.78
0.77
0.75
0.77
CO+H2O
(b)
CO2+H2
very slow
fast
HCOOH
CO2+H2
very slow
fast
HCOOH
CO+H2O
Figure 6. Control of the WGS reaction to produce
formic acid.
xHCOOH
t/h
0
0.33
3.2
6.0
9.0
12.0
15.0
19.7
Table 3. The time evolution of the reactant and
product yields for the decomposition of 1.0 mol kg-1
formic acid at 330 °C and ρT ≈ 0.99 (the volume of
the gas phase is negligible). The temperature is too
high to follow the decarbonylation kinetics.
1.0
ρ T = 0.30
0.8
0.50
0.80
0.6
0.4
0.90
0.2
1.00 (totally filled)
0.0
180
200
220 240
o
T/ C
260
280
Figure 7. The temperature dependence of
formic acid yield xHOOH. Each of the contours
corresponds to a fixed filling factor ρT
indicated in the figure.
8
Another possible source of the formic acid fuel
is the electric energy from wind and solar power
plants. The electricity is first used to electrolysis
water into hydrogen and oxygen. The hydrogen is
then converted into formic acid which can be
conveniently stored. The conversion of hydrogen
into formic acid accompanied by carbon dioxide
fixation is the reverse reaction of the
decarboxylation of formic acid. To convert
hydrogen into formic acid, the reaction conditions
should be chosen so that the decarbonyaltion rate of
formic acid is much slower than that of the reverse
decarboxylation; see Fugure 6b. Such conditions
would be satisfied in the presence of such a weak
base as sodium hydrogencarbonate (note: H2 +
NaHCO3
HCOONa + H2O) or metal catalyst
such as SUS316 or more powerful but expensive
rhodium complexes. The equilibrium formic-acid
yield from hydrogen and carbon dioxide is
determined by the value of the equilibrium constant
KCO2. In basic conditions, the neutralization Gibbs
energy change for formic acid and carbon dioxide
should be also taken into account. Although the
KCO2 value determined in acidic conditions implies
an unfavorable formation of formic acid, a higher
formic acid yield can be expected in basic
conditions [10] and at a lower temperature [5d] in
the presence of metal catalysts.
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In this paper we have established the basic
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9
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10