Extracto
1
The Eco-Friendly and Promoting Influence of
Nitric AcidSteam Vapors on the Oxidation of
C
3
C
4
Parrafins into Methanol
AYODEJI A. IJAGBUJI
2
Abstract
The catalytic conversion of light alkane feedstock is one of the promising industrial routes to prepare
valuable chemical products: owing to its use in car and gas turbine engines, heaters, incinerators, and
hydrocracking furnaces. Although the combustion of hydrocarbon fuels appears to be conceptually
simple; the details of how alkane conversion to carbon dioxide and water occurs with concomitant
release of energy, are enormously complex. Without timely and effective measures to mitigate the
adverse impacts of greenhouse gases emission, the living environment of the world will continue to
deteriorate and become increasingly unbearable. On that ground, the development of technological
methods for methanol production via alkane conversion in a manner that is energy-conserving,
inexpensive, environmentally-sustainable, and least-damaging to human health & welfare are of
crucial importance. The applied method in this work contributes to these objectives.
It is well known that in the petroleum refining process, a substantial amount of C
3
C
4
fraction
(existing as a mixture rather than individual component), is primarily recovered from crude oil
distillation and by cracking of heavy molecules: bulk of which is flared as hydrocarbon fuel into the
atmosphere through refinery furnace. On that ground, owing to the higher chemical reactivities of
propane and butane, in comparison to methane, the use of C
3
C
4
hydrocarbon mixtures has been
investigated as an alternative feedstock for methanol production, which apparently, to a large extent,
is an important aspect of chemical technology and economics.
The present work evaluates the theoretical concepts, quantum chemical calculations, and
experimental investigations to explore novel pathways for the direct oxidation of propane and butane
fraction to methanol using a photochemical reactor.
To test the validity of the mechanism, the photochemical oxidation of C
3
C
4
gas fraction was
investigated under three experimental phases under mild reaction conditions at a temperature (T =
100°C), at a pressure (P = 1 atm.), and at reaction times (t
r
) within 5 120mins (a) without exposure
to irradiation, (b) with exposure to visible light irradiation at a wavelength region ( = 420 nm), (c)
with exposure to visible light irradiation at a wavelength region ( = 420 nm) under the auto-catalytic
influence of nitric acid vapor. The oxidation products obtained are methanol, as well as ethylene and
propylene (very important starting materials in the petrochemical industry).
This proposed methodology, which is initiated with a recently developed technology specifically to
regenerate nitric acid and significantly minimize energy waste: can be considered as a useful guidance
to design an industrial plan for the direct conversion of C
3
C
4
fraction into methanol.
3
Contents
PREFACE
ii
ACKNOWLEDGEMENTS
iii
CONTENTS
iv
LIST OF FIGURES
vii
LIST OF TABLES
xi
1. INTRODUCTION...1
1.1 Technical background... 2
1.1.1 Methanol synthesis: A brief history...2
1.2 Production of methanol: From syn-gas, bio-sources, methyl formate, & CO
2
recycling...4
1.2.1 Methanol from biomass and cellulosic sources...4
1.2.2 Methanol from chemical recycling of CO
2
...6
1.2.3 Methanol from methyl formate...8
1.2.4 Methanol from syn-gas...10
1.2.4.1 Syn-gas preparation techniques...11
1.2.4.1.1 Syn-gas via steam reforming...12
1.2.4.1.1.1 Steam reforming of methane (SMR)...12
1.2.4.1.1.2 Catalytic partial oxidation of methane (CPO)...13
1.2.4.1.1.3 Auto-thermal reforming (ATR)...13
1.2.4.1.1.4 CO
2
reforming of methane...14
1.2.4.1.2 Syn-gas from petroleum oil & higher hydrocarbons...14
1.2.4.1.3 Syn-gas from coal...15
1.2.4.1.4 Syn-gas via combined-reforming...15
1.2.4.1.5 Syn-gas via heat-exchange reforming...16
1.3 Methanol: Application & Economy...16
1.3.1 Methanol as a chemical feedstock...16
1.3.2 Methanol as transportation fuel...18
1.3.2.1 Methanol as fuel in internal combustion engines (ICE)...18
1.3.2.2 Methanol as fuel in compression ignition engines (DIESEL)...19
1.3.3 Methanol as gasoline additive...19
1.3.4 Methanol for static power & heat generation...20
4
1.3.5 Methanol for waste-water denitrification...20
1.3.6 Methanol for bio-diesel trans-esterification...20
1.4 Market dynamics for methanol...21
2. LITERATURE REVIEW...23
2.1 Present-day investigations...23
2.2 Main products of the DMTM process...24
2.2.1 The CH
4
/O
2
ratio...25
2.2.2 The CH
3
OH/CH
2
O ratio...26
2.2.3 The CO/CO
2
ratio...28
2.2.4 By-products of the DMTM process...28
2.2.5 Yield of methanol & oxygenates...29
2.3 Main parameters of the DMTM Process...32
2.3.1 Effect of oxidant on the selectivity & yield of products...32
2.3.1.1 Influence of oxygen concentration on the temperature and rate of reaction...33
2.3.2 Effect of temperature on the yield of products...36
2.3.3 Effect of pressure...37
2.3.3.1 Effect of pressure on the temperature and rate of reaction...39
2.3.3.2 Effect of pressure on the rate of branched-chain quasi-stationary reaction...40
2.3.4 Effect of gas composition (3
rd
body effect)...42
2.3.4.1 Heavier homologues of methane...42
2.3.4.2 Inerts (N
2
, He) ...44
2.3.4.3 Carbon oxides (CO, CO
2
)...45
2.3.5 Effect of homogeneous promoters...47
2.3.6 Effect of heterogeneous catalysts...57
2.3.6.1 Effect of catalyst's surface area...60
2.3.7 Effect of feed-flow rate (residence time)...60
2.3.8 Surface-effect of reactor material...62
2.3.8.1 Role of diffusion of reactants to the reactor surface...65
2.3.8.2 Decomposition of products on the reactor surface...67
2.4 Key features of the mechanism...69
2.4.1 Mechanism for the gas-phase oxidation of methane in medium temperature range...69
2.4.2 Main kinetic features of the DMTM process...72
5
2.5 Role of heterogeneous process in the DMTM process...77
2.5.1 The interplay between the homogeneous and the heterogeneous catalytic processes of methane
oxidation...78
2.6 Thermo-kinetic phenomena in partial oxidation of methane...83
2.6.1 Experimentally-observed thermo-kinetic phenomena in the partial oxidation of methane...83
2.7 Effect of physical promotion of the process...88
2.8 Effect of the process under supercritical conditions...94
2.9 Overview of experimental achievements on DMTM process...98
2.10 Conclusions...99
3. AIM & JUSTIFICATION FOR THE RESEARCH DIRECTION...100
3.1 Research justification...100
3.2 Aim of research...101
3.3 Theoretical concepts & mechanism for the autocatalytic photo-oxidation process of
propane-butane fraction to methanol...101
4. EXPERIMENTALS...105
4.1 Experimental set-up...105
4.2 Experimental procedures...106
5. RESULTS AND DISCUSSION...108
6. CONCLUSIONS...119
REFERENCES...121
6
LIST OF FIGURES
Fig. 1.1
Scheme for methanol production via biomass...6
Fig. 1.2
Scheme for methanol production via natural gas...11
Fig. 1.3
Methanol derived chemical products & materials...17
Fig. 2.1 (a)
Selectivity of CH
3
OH (solid lines) and CH
2
O (dotted lines) formation at P = 1.5(),
2.0(), 3.0 atm(),
(b)
Pressure dependence of the CH
3
OH/CH
2
O ratio; on CH
4
conversion...27
Fig. 2.2
Pressure dependence of concentrations of carbon oxides and [CO]/[CO
2
] ratio at T = 400
o
C
and [O
2
]
2.8%: (1) CO, (2) CO
2
, and (3) [CO]/[CO
2
]...28
Fig. 2.3
Dependence of CH
3
OH:
(a)
Selectivity and
(b)
Yield; on CH
4
conversion obtained from fast-
flow experiments at t
r
< 20s and P 30 atm...29
Fig. 2.4
Dependence of CH
3
OH:
(a)
Selectivity and
(b)
Yield; on CH
4
conversion obtained from
slow-flow experiments at t
r
> 20s and P 25 atm...30
Fig. 2.5
Dependence of the yield of liquid oxidation products on oxygen concentration...31
Fig. 2.6
Dependence of the yield of liquid products on CH
4
/O
2
ratio at P = 100 atm and T = 400
o
C: (1)
Total liquid products yield, (2) water, (3) methanol, and (4) formaldehyde...33
Fig. 2.7
Dependence of: (
a
) The yields Y (g/m
3
) of sum of liquid products (1, 2), methanol (3, 4), and
formaldehyde (5, 6); (
b
) Concentration of water (1, 2), methanol (3, 4), and formaldehyde (5, 6): on
initial oxygen concentration in quartz (1, 3, 5) and (2, 4, 6) stainless steel reactor at T = 550
o
C, P = 80
atm...34
Fig. 2.8
(a)
Temperature for the ignition of CH
4
/O
2
mixture at (1) 40%, (2) 29%, (3) 21%, (4) 10% O
2
conversion at P = 103 atm,
(b)
Dependence of time of (1) 5%, (2) 50%, (3) 95% O
2
conversion at T =
693K, P = 10 atm...34
Fig. 2.9
Reactor temperature profiles at P = 50 atm., initial temperatures, and various oxygen
concentrations...35
Fig. 2.10
Methanol selectivity as a function of reactor temperature at different oxygen concentrations;
TF is the total flow of the reactants, cm
3
/min...36
Fig 2.11
Temperature profiles along tubular reactor for various preset reactor wall temperatures at a
flow rate of 1000 ml/min, P = 3.0MPa, and [O
2
]
0
= 5%...37
Fig. 2.12
Dependence of the methanol yield on pressure. (I) calculated by kinetic model at T = 480
o
C,
[O
2
] = 3.5%, and a reactor diameter of 10mm, (II) calculated by the equation [CH
3
OH] = [CH
3
OH]
/(1 + P
0.5
/P)...38
7
Fig. 2.13
Dependence of (
a
) yields Y (g/m
3
) of (1) water, (2, 3) methanol, and (4) formaldehyde in a
(1, 2, 4) quartz and (3) stainless steel reactor; (
b
) The concentration (1) water, (2) methanol, and (3)
formaldehyde in liquid products in quartz reactor ([O
2
]
0
= 3.6% and T =550
o
C) on pressure...38
Fig. 2.14
Dependence of the selectivity and products ratio on pressure () S
COx
, () S
C2+
, () CO
x
/C
2+
,
() C
2
H
6
/C
2
H
4
...39
Fig. 2.15
Dependence of the (
a
) temperature of complete oxygen conversion at [O
2
]
0
= 14% and
reaction time t
r
= 0.6 0.8s, (
b
) reaction time t
r
at T = 650K on pressure...40
Fig. 2.16
Effect of addition of 3% C
2
-C
4
hydrocarbons on:
(a)
CH
4
conversion;
(b)
CH
3
OH yield; at
P = 41 atm, t
r
= 1s, CH
4
: O
2
: N
2
= 30 : 1 : 4...41
Fig. 2.17
Dependence of
(a)
CH
4
conversion,
(b)
CH
3
OH yield; on oxygen concentration for the
oxidation with () oxygen, and () air at P = 40 atm, T = 440 460
o
C, and t
r
= 0.6 1.0s...45
Fig. 2.18
Dependence of yield of liquid oxidation products on CO concentration at P = 100 atm and T
= 400
o
C: (1) Total yield of liquid products, (2) water, (3) methanol, and (4) formaldehyde...46
Fig. 2.19
The production of formaldehyde in empty quartz reactor: P = (- - -) 1, and (...) 5 atm., (-.-.)
in the presence of NO (empty reactor), and in the presence of different catalysts...49
Fig. 2.20
Effect of NO addition on light alkanes conversion: CH
4
(,); C
2
H
6
(,); and C
3
H
8
(,):
Filled and unfilled symbols are experiments with and without NO additives, respectively...51
Fig. 2.21
Selectivity of product formation as a function of NO
2
concentration at 10% CH
4
conversion
(55.6% CH
4
, 27.7% O
2
, NO
2
and He; mixture flow rate, 120 ml/min; P = 1 atm: CO (); CH
3
OH ();
HCHO (); CH
3
NO
2
(); and CO
2
()...53
Fig. 2.22
Main conversion pathways of intermediate products for CH
4
oxidation in the presence of
NO
x
...56
Fig. 2.23
Comparison of CH
4
conversion and selectivity of CH
3
OH formation in homogeneous and
catalytic reactions: P = 50 atm, [O
2
] = 4.35%, and t
r
= 2s in an empty reactor and 4.5s in the presence
of the catalyst. The reactor was filled with pyrex spheres (P), V
2
O
5
, and Ag/-alumina...58
Fig. 2.24
Dependence of CH
3
OH yield on the [CH
4
]/[O
2
] ratio in the absence (), and presence () of
Cu/ZnO/Al
2
O
3
catalyst at P = 4 atm, and T = 575
o
C...59
Fig 2.25 (a)
Product selectivity as a function of residence time at T = 400
o
C, P = 31 atm., CH
4
-O
2
-N
2
= 30 : 1 : 4,
(b)
Time evolution for the yield of various products in Cu reactor at P = 300 atm, and T =
375
o
C: CH
3
OH (), CO (), CH
2
O (), CH
2
O
2
()...61
8
Fig. 2.26
(a) Kinetic curves at long induction period at P = 100 atm, [O
2
]
0
= 3%, T = 755K, (b)
Dependence of oxygen conversion (
) and the yields of CH
3
OH (), and CH
2
O () on the reaction
time at [CH
4
]/[O
2
] = 16, T = 400
o
C, P = 34 atm...62
Fig. 2.27 (a)
CH
4
conversion and,
(b)
CH
3
OH selectivity; versus reactor wall temperature in the
presence (o) and absence () of quartz filling spheres at a gas flow rate of 200 ml/min, P = 3.0MPa,
and [O
2
]
o
= 5%...64
Fig 2.28
Temperature of the degree of decomposition of CH
3
OH on the surface of copper (),
stainless steel (), and pyrex () at P = 30 - 50 atm., [O
2
]
o
= 2.5 - 4.9%, t
r
= 200s...67
Fig. 2.29
Effect of reactor diameter on the selectivity of CH
3
OH formation at P = 91 atm, T = 427
o
C,
t
r
= 35s: d = 11mm () ; d = 19mm (); d = 30mm () ; and d = 30mm ()...69
Fig. 2.30
Kinetics of the formation of CH
3
O
2
radicals at T = 650K and P =
(a)
3;
(b)
4 atm...72
Fig. 2.31
Calculated time histories of CH
3
O
2
radical concentration for:
(a)
The process as a whole;
(b)
The 1
st
stage and 2
nd
stage at PCH
4
= 82 atm, T = 406
o
C and PO
2
= (1) 1.8 atm, (2) 8.5 atm...74
Fig. 2.32
Kinetics of the formation of partial methane oxidation products at T = 783K, P = 85 atm.,
CH
4
:O
2
= 21:1; (1) CH
2
O, (2) H
2
O, (3) H
2
O
2
, (4) CH
3
OH, (5) CO, (6) O
2
, (7) CO
2
...75
Fig 2.33
Dependence of
(a)
calculated CH
3
OH yield and reaction time t
r
on ratio of (W eff/W therm)
at T = 410
o
C, P = 100 atm, [CH
4
]/[O
2
] = 19,
(b)
CH
2
O and reaction time t
r
on the ratio of log (W
eff/W therm) at T = 753K, P = 1 atm, [CH
4
]/[O
2
] = 9:1...81
Fig 2.34
Dependence of the (, o) CH
4
conversion and (, ) CH
3
OH yield for the rising (solid line)
and lowering (dashed line) temperature of reactor walls at P = 30 atm, [O
2
] = 9.5%, t
r
= 0.2s...83
Fig 2.35
Simulation of self-heating of the reaction (T) and CH
3
OH concentration on T
wall
at [O
2
] =
9.5%, P = 50 atm, and t
r
= 10s. The arrows show temperature variation...84
Fig 2.36
Simulation of methanol formation selectivity on the temperature of the reactor walls T
wall
at
P = 30 atm and [O
2
] = () 9.5%, () 7.5%, () 5%, and () 2.5%...85
Fig. 2.37
Kinetic curves of (a) heating and (b) emission for the oxidation of a 2CH
4
-O
2
mixture at T =
784K and various pressures: P (Torr) = (1) 585, (2) 650, (3) 683, (4) 709, and (5) 727...86
Fig. 2.38
Cool-flame region for the oxidation of 2CH
4
-O
2
mixture: (1) single- and (2) double-cool
flame flashes...87
Fig. 2.39
Versions of design of the quartz plasma reactor...91
Fig. 2.40
Influence of specific energy input on (a) CH
4
conversion and (b) CH
3
OH yield at P = 2 atm.,
T
wall
= 80
o
C, CH
4
/O
2
= 8 : 2, and t
r
= 6s. The power was varied at a flow rate of 1.0NL/min; the flow
rate was varied at a power of 200W...92
9
Fig. 2.41 (a)
Temperature dependence of the rate constant k,
(b)
Selectivity dependence of CH
3
OH
formation; on CH
4
/O
2
ratio (70% H
2
O-30% CH
4
) in SCW...97
Fig. 4.1
Schematic diagram of the pilot unit...105
Fig. 5.1
Quantum chemical results of the B3LYP/6-311++G(3df,3pd) calculations for the optimized
structures: (a) Molecular adduct of H
2
O and NO
2
*
; (b) Transition state (TS) for reaction: NO
2
*
+ H
2
O
HONO +
·
O...110
Fig. 5.2
Calculated relative energy diagram for the photochemical reaction: NO
2
*
+ H
2
O HONO
+
·
OH...112
Fig. 5.3
Quantum chemical results of the B3LYP/6-311++G(3df,3pd) calculations: (a) Optimized
structure of molecular adduct H
2
ONO
2
; (b) Transition state (TS) for reaction of methyl radical with
water molecule...114
Fig. 5.4
Spectral response of the light source (lamp DRT 100)...118
10
LIST OF TABLES
Table 1.1
Physicochemical parameters for methanol...3
Table 2.1
Gibbs free energies of reactions at different temperature...25
Table 2.2
References and experimental conditions for Fig. 2.28...41
Table 2.3
Kinetic parameters for the reaction of C
1
C
4
alkanes with oxygen...44
Table 2.4
Activation energy of the gas-phase oxidative conversion of light alkanes in the presence of
NO additives...50
Table 2.5
Products of gas-phase reaction of CH
4
-O
2
-NO
x
system under the same conditions...53
Table 2.6
Mechanism of the initial stage of the DMTM Process...71
Table 2.7
Thermodynamic and critical parameters for C
1
C
4
alkanes...95
Table 2.8
Experimental conditions and main products of CH
4
oxidation to CH
3
OH in supercritical
water...96
Table 3.1
Physicochemical parameters of natural gas, propane & butane...100
able 5.1
Structural parameters (R
NO
, ), spin density on atoms
s
, dipole moments (D), vibration
frequency (m
-1
) and 0-0 transition energy (T
0
, eV) for NO
2
molecule in the ground (
2
A
1
) and the
excited states (
2
B
2
,
2
B
1
,
2
A",
2
A
2
) B3LYP/6-311++G(3df,3pd) approach...108
Table 5.2
Total energy (E
p tot
, a.u.), vertical transitions ( E
op
, eV), transition moments (M
op
, Debye),
polarization of the light () and oscillator strength (f
op
) for transitions from the ground state NO
2
calculated by the MCSCF-CI/6-311++G
*
. A
p
is the Einstein Coefficient (10
4
s
-1
) for spontaneous
emission to the ground state...110
Table 5.3
Results of B3LYP/6-311++G(3df,3pd) calculations for the total energy (E
total
), zero-point
vibration energy (E
0
), and absolute entropy (S
o
298
) for the reactants NO
2
(
2
A") and H
2
O (
1
A
1
), their
molecular adduct, transition (TS) state, and thermodynamic parameters
r
G
0
298
,
r
H
0
298
and
r
S
0
298
for
the reaction: ON-O
·
+ H
2
O ONO +
·
O...113
Table 5.4
Results of quantum chemical calculation for total Energy (E
total
), Zero vibrational energy
(E
0
), and absolute entropy S
0
298
for ground and excited-state molecule and the thermodynamic
parameters (
r
G
0
298
,
f
H
0
298
,
r
S
0
298
) for reaction:
·
CH
3
+ (H
2
ONO
2
) CH
3
OH + HNO
2
...115
Table 5.5
Results of quantum chemical calculation for total Energy (E
total
), Zero vibrational energy
(E
0
), and absolute entropy S
0
298
for ground and excited-state molecule and the thermodynamic
parameters (
r
G
0
298
,
r
H
0
298
,
r
S
0
298
) for reaction:
·
CH
3
+ (H
2
ONO
2
) CH
3
NO
2
+ H
2
O...116
Table 5.6
Effect of the Experimental Conditions on the Yield of Methanol...117
11
Chapter 1
INTRODUCTION
Methanol
is an oxygenated hydrocarbon fuel, which also, serves as a raw material for chemical
synthesis in the petrochemical industry. Thereby, it is considered as one of the most-traded industrial
chemicals, and an integral part of our economy
[1]
(with a production capacity of approximately 20
billion gallon annually).
Conventional technologies for methanol production involves steam methane reforming (SMR) to
syn-gas at high temperature (800 1000
o
C) and moderate pressure (4Mpa) in Eq. (1.1), followed by
syn-gas conversion to methanol in Eq. (1.2) by passing it over co-precipitated copper/zinc-oxide
(Cu/ZnO) catalysts at low temperature (220 280
o
C) and high pressure (5 8MPa).
[2]
CH
4
+ H
2
O CO + 3H
2
H
0
298
= + 206kJ/mol
(1.1)
CO + 2H
2
CH
3
OH
H
0
298
= 90.7kJ/mol
(1.2)
However, the overall efficiency of methanol synthesis is greatly undermined by thermodynamics, and
a high-theoretical carbon monoxide (CO) conversion (20% at 280
o
C) due to the limitations imposed
by the reaction equilibrium and low heat efficiency. In addition, the existing commercial process for
methanol production via SMR is very energy-intensive, exhaustive, and not cost-effective.
[3]
To significantly minimize the methanol production capital cost (inherent in the steam-reforming
stage), an economically-feasible one-step methane conversion into methanol, is an attractive option.
The direct partial oxidation of methane (POM) into methanol Eq. (1.3) is an exothermal reaction that
is energetically more-efficient, tends to minimize temperature-and-pressure fluctuations, decreases
the second-law efficiency losses, and can give quite good results than the endothermic steam-methane
reforming (SMR) reaction.
[4]
Nevertheless, it is operated by a free radical mechanism which is hard to
control.
CH
4
+ ½O
2
CH
3
OH
H
0
298
= 126.4kJ/mol
(1.3)
Also, the rapid subsequent combustion reactions of methanol to form CO
2
, limit conversions in
practice, while the explosive nature of the required reactant mixtures creates engineering challenges in
the design of mixing schemes and pressure vessels.
12
1.1
TECHNICAL BACKGROUND
1.1.1
METHANOL SYNTHESIS A BRIEF HISTORY
Methanol in its relatively pure form was first isolated in 1661 by Robert Boyle
[5]
who called it spirit of
the box because he produced it through the distillation of boxwood. Its chemical identity or elemental
composition, CH
3
OH, was described in 1834 by Jean-Baptiste Dumas and Eugene Peligot. They also
introduced the word methylene to organic chemistry, from the Greek words "methu and hyle",
meaning respectively wine and wood. The synthetic route to methanol production, by reacting carbon
monoxide with hydrogen, was first suggested in 1905 by the French chemist Paul Sabatier.
[6]
In 1913, the Badische Anilin und Soda Fabrik (BASF), based on the investigations of A. Mittasch
and C. Schneider, patented a process to synthesize methanol from synthesis gas, produced from coal,
over a Z
n
/C
r
O catalyst at 300 400
o
C and 250 350 atmospheres.
[7]
Between 1923 and 1926, F. Fischer and H. Tropsch reported from the Muhlheim Coal Research
Laboratory their extensive studies of the production of hydrocarbons, including that of methanol from
syn-gas, a mixture of carbon monoxide and hydrogen, which was the basis of what is known as the
FischerTropsch synthesis.
[8 10]
Methanol production in 1927, started in the U.S. by the Commercial Solvents Corporation and the
DuPont. The Commercial Solvents Corporation process used the high-pressure technology to produce
methanol from CO
2
/H
2
mixtures obtained as fermentation by-product gases
[11]
, whereas, DuPont
process utilized coal to produce gaseous feedstock (synthesis gas).
The steam reforming process of methane, because of the very high purity of the synthesis gas,
opened the way to the technical realization of low-pressure methanol process, introduced
commercially in 1966 by Imperial Chemical Industries (ICI). This new process, using a more active
C
u
/Z
n
O catalyst and operating at 250 300
o
C and 100 atm.,
[12]
puts an end to the high-pressure
methanol synthesis technology, which operated under much more severe conditions. Notwithstanding,
these active C
u
/Z
n
O catalysts were highly susceptible to poisoning and deactivation (from
overheating), thus making close control of methanol reactor extremely important.
[13]
Advances in synthesis gas production and purification removed most of the catalyst poisons and
made use of the new catalysts possible. These low-pressure methanol synthesis catalysts have a life
time of up to 4 years. Further advancements have led to the invention of many of such catalysts and
process conditions for effective production methods.
Today Worldwide, methanol is produced from synthesis gas. However, current research is aimed at
developing new ways to synthesize methanol at even lower temperature and pressure via:
13
hydrogenative chemical recycling of carbon dioxide, as well as, directly from methane (natural gas),
without prior formation of synthesis gas (a superior approach from an energetic viewpoint), are being
developed. The physicochemical properties of methanol is shown in Table 1.1.
TABLE 1.1
PHYSICOCHEMICAL PARAMETERS FOR METHANOL
Synonyms
Methyl alcohol, Wood alcohol
Chemical formula
CH
3
OH
Molecular weight
32.04 g/mol
Chemical compositions:
Carbon (C)
37.5%
Hydrogen (H)
12.5%
Oxygen (O)
50%
Critical temperature
512.5K (239
o
C, 463
o
F)
Critical pressure
8.084Mpa (78.5 atm)
Critical density
0.2715g cm
3
Critical compressibility factor
0.224
Specific gravity @ 25
o
C
0.7866
Melting point
- 97.6
o
C
Boiling point
64.6
o
C
Reid vapor pressure flash point
32kPa
Density @ 20
o
C
791kgm
-3
Energy content
5420kcal kg
-1
(173.6 kcal mol
-1
)
Latent heat of vaporization
9.2kcal mol
-1
Flash point
11
o
C
Auto-ignition temperature
455
o
C
Surface tension @ 25
o
C
22.07mN m
1
Refractive index
1.32652
Heat of combustion @ 25
o
C, 101.325Kpa:
Higher heating value (HHV)
22.7kJg
1
Lower heating value (LHV)
19.9kJg
1
Thermal conductivity @ 25
o
C
200mWm
1
K
1
Coefficient of cubic thermal expansion @ 20
o
C
0.00149 per
o
C
Auto-ignition temperature viscosity @ 25
o
C
0.544mPa s
14
Limits of flammability (in air):
Lower
6.0 % (v/v)
Higher
36.5 % (v/v)
1.2
PRODUCTION OF METHANOL: FROM FOSSIL FUELS, BIO-SOURCES, AND
CHEMICAL CARBON-DIOXIDE RECYCLING
1.2.1
METHANOL FROM BIOMASS AND CELLULOSIC SOURCES
Biomass is referred to as any type of plant or animal material i.e., materials produced by life forms.
This includes high-cellulose content materials, whole plant matter, industrial-agricultural residue,
wastes (municipal solid waste, animal waste), lignite, anaerobic digestion, and domestic effluent from
sewage treatment plants etc.
Methanol as mentioned was originally made from thermal destructive distillation of wood, but due
to its bulkiness, low-energy density, relatively higher energy intensity, lower yield (compared to
natural gas), greater logistics (solids handling versus pipeline), and the advent of synthesis via syn-gas,
this route was rapidly abandoned in the first half of the 21
st
century. This process involves the
gasification of biomass to syngas, followed by the synthesis of methanol.
In the gasification process, the biomass feedstock is usually first dried and pulverized to yield
particles of a uniform size and with moisture content no higher than 15 20% for optimal results.
Gasification is a thermochemical process that converts biomass at high temperature into a gas
mixture containing hydrogen, carbon monoxide, carbon dioxide, water vapor and light hydrocarbon.
The pretreated biomass is sent to the gasifier where it is mixed, generally under pressure, with oxygen
and water. Gasification of biomass feed-stocks is therefore generally a two step process:
In the 1
st
stage of gasification pyrolysis or destructive distillation; the dried biomass is heated to
temperatures ranging from 400 600
o
C in an atmosphere too deficient in oxygen to allow complete
combustion. The pyrolysis gas obtained consists of carbon monoxide, hydrogen, methane, volatile tars,
carbon dioxide and water. The residue, which is about 10 25% of the original fuel mass, is charcoal.
In the 2
nd
stage of gasification char conversion; charcoal residue from the pyrolysis step reacts with
oxygen at 1300 1500
o
C, producing carbon monoxide. The syn-gas obtained has to be purified
before its transformation via catalysis to methanol at a temperature of 250
o
C and pressure range of
35 90 bar.
[14]
A recent report revealed that methanol can be produced from biomass with a net higher
heating value (HHV) energy efficiency between 50 57%,
[15]
and an HHV value for methanol = 22.9
MJ/kg.
[16]
15
Since the transportation of biomass products over long distances is not economical, its transformation
into an easy to handle-and-store liquid intermediate through fast pyrolysis, has, recently been
proposed. In fast pyrolysis, small biomass particles are heated very rapidly to 400 600
o
C at
atmospheric pressure to yield oxygenated hydrocarbon gases. The generated gases are then rapidly
quenched to avoid their decomposition by cracking. The black liquid resembling crude oil is obtained
called "bio-crude."
[17]
Biomass gasification has a comparative advantage over coal such as its much lower sulfur content
(0.05 0.20%wt), while heavy metals impurities (mercury, arsenic, etc.), are present only in
insignificant quantities, however, it is: (i) hugely, energy-intensive, (ii) expensive biomass
gasification resulting to an overall bio-methanol production costs (1.5 4 times higher than natural
gas), (iii) formation of tars, dust and inorganic substances, in particular, is most cumbersome and
problematic for the commercialization of any biomass gasification technology.
[18]
These tars:
composed mainly of oxygenated compounds and higher molecular weight hydrocarbons are
problematic because they condense in pipes, boilers, transfer lines and particulate filters, leading to
blockage, clogging and other operational difficulties, and (iv) low carbon conversion due to the
stoichiometric adjustment and the high purge gas rate. To achieve high carbon conversions for
methanol synthesis from biomass feed-stocks, external hydrogen admixture to the gasifier syngas is
required (HYNOL PROCESS).
Fig. 1.1
A Scheme for Methanol Production via Biomass.
[19]
16
A train of processes to convert biomass to required gas specifications precedes the methanol reactor.
These processes include pre-treatment, gasification, tar cracking, wet-gas cleaning, gas-conditioning,
reforming of high hydrocarbons, water-gas-shift reaction (WGSR), H
2
addition and/or CO
2
removal
(to obtain an optimized syn-gas), methanol synthesis, and purification,
[19]
as depicted in
Fig 1.1
1.2.2
METHANOL PRODUCTION VIA CHEMICAL RECYCLING OF CO
2
Atmospheric carbon dioxide is the most abundant greenhouse gas, and the scientific community now
agrees that its increasing concentration in the atmosphere is mainly anthropogenic in origin. The use
of carbon dioxide (CO
2
) as a raw material for chemical synthesis is the subject of growing attention,
driven by environmental, legal, and social concerns.
[20]
As an economical, safe, and renewable carbon source, CO
2
turns out to be an attractive C
1
building
block for making organic chemicals. Due to the high kinetics and thermodynamic stability of CO
2
,
high energy substances or electro-reductive processes are typically required to transform CO
2
into
other chemicals.
[21 22]
As an alternative feedstock, CO
2
has replaced CO, and is considered as an
effective approach for methanol production.
CO
2
+ 3H
2
CH
3
OH + H
2
O,
H
0
298
= 49.5 kcal/mol
(1.4)
O=C=O + H
2
HOCH=O
(1.5)
HOCH=O + H
2
HOCH
2
OH CH
2
=O + H
2
O
(1.6)
CH
2
=O + H
2
CH
3
OH
(1.7)
From the thermodynamic viewpoint, a decrease in reaction temperature or an increase in reaction
pressure could favor the synthesis of methanol. Indeed, enhanced reaction temperature > 513K
facilitates CO
2
activation and subsequent methanol formation. Furthermore, other by-products such as
CO, hydrocarbons, and higher alcohols are formed during CO
2
hydrogenation.
Therefore, to avoid the formation of these undesired by-products, a highly-selective catalyst is
required. Both homogeneous and heterogeneous catalysts play crucial roles in CO
2
hydrogenation.
Homogeneous catalysts (e.g. Ru, Rh, and Ir) show satisfactory activity and selectivity, but the
recovery and regeneration are problematic. Alternatively, heterogeneous catalysts (e.g. Fe, Cu, Ni) are
preferable in terms of stability, separation, handling, and reuse, as well as reactor design, which
reflects in lower costs for large-scale productions. However, these heterogeneous catalysts frequently
suffer from low yield and poor selectivity due to fast kinetics of the CH bond formation. Thereby, the
use of expensive catalysts,, high temperature, high operating pressure, and the tedious workup
procedures involved for catalyst separation and recycling make these processes unattractive for
17
commercial applications.
In order to make CO
2
hydrogenation economically feasible, significant improvements such as the
exploitation of novel heterogeneous catalysts,
[23]
immobilization of homogeneous catalysts,
[24 25, 29]
and the use of green solvents such as ionic liquids (IL's) and supercritical CO
2
(scCO
2
)
[26 28]
in new
catalytic systems with rational designs and molecular simulations are required.
Even though a large number of investigations have been done with experimental observations and
theoretical analyses using varied metal-based catalysts and modifiers, yet, the mechanisms for CO
2
hydrogenation are still in dispute: for instance, fundamental understanding regarding the role of added
solvent in homogeneous systems (at the molecular level) is unclear. In heterogeneous reaction, the
prevalent consensus is that the active site is provided by the synergy between the primary catalyst and
the support or the promoter.
Nevertheless, the nature of the active sites and interactions among active components, support, and
promoter as well as reaction mechanisms are still elusive, i.e. for both homogeneous and
heterogeneous catalysts, the primary focus of the mechanisms for CO
2
hydrogenation is on how and
where CO
2
is activated and interacts with hydrogen and/or hydroxyl species under different reaction
conditions.
Industrial utilizations of CO
2
as solvent and reactant amount to only 0.5 wt% (~ 128Mt/y) of the total
anthropogenic CO
2
emissions every year. In principle, chemical utilizations of CO
2
do not necessarily
help mitigate the greenhouse effect considering energy input and carbon circulation.
However, if CO
2
could be chemically transformed to fuels, it would be helpful to circulate carbon to
alleviate the greenhouse effect. Particularly, production of fuels that can be easily stored and
transported is preferable. The capital investment for a methanol plant using CO
2
and H
2
is estimated
to be about the same as that for a conventional syn-gas based plant.
[30]
The limiting factor for large
scale-up of such processes is the availability and price of CO
2
and H
2
and, first of all, the necessary
energy.
From a scientific standpoint, the capture of CO
2
from the atmosphere, and its recycling from
industrial or natural emissions would provide a renewable and inexhaustible carbon source, allow the
continued use of derived carbon fuels in an environmentally carbon neutral way, and offer
opportunities to mitigate the increasing CO
2
buildup. The hydrogenation of chemically stable CO
2
to
methanol is a highly attractive solution to the recycling of the carbon source and has favorable
thermodynamics, but high activation energy barriers should be overcome by appropriate catalysts.
18
1.2.3
METHANOL PRODUCTION FROM METHYL FORMATE
To reduce the pressure and temperature needed for the current methanol production process, and also
to improve its thermodynamic efficiency, alternative routes to convert CO/H
2
mixtures into methanol
under milder conditions have been developed.
Among these, the most notable is the synthesis of methanol via methyl formate, first proposed in
1919 by Christiansen.
[31]
CH
3
OH + CO HCOOCH
3
(1.8)
HCOOCH
3
+ 2H
2
2CH
3
OH
(1.9)
CO + 2H
2
CH
3
OH
This methanol synthesis route consists of two steps: Methanol is first carbonylated to methyl formate
in Eq. (1.8), which is subsequently reacted with hydrogen to produce two moles of methanol in Eq.
(1.9). The carbonylation reaction is carried out in the liquid phase using sodium or potassium
methoxide (NaOCH
3
or KOCH
3
) as a homogeneous catalyst. This is a proven and commercially
available technology used in the production of formic acid from CO through methyl formate. High
activities for methanol carbonylation have also been recently shown with Amberlyst and Amberlite
resins used as heterogeneous catalysts.
The subsequent reaction of methyl formate with hydrogen (hydrogenolysis) to produce methanol
can be conducted either in the liquid or gas phase using a copper-based catalyst (copper chromite,
copper supported on silica, alumina, magnesium oxide, etc). Carbonylation and hydrogenolysis can be
carried out in two separate reactors, but are preferably combined in a single reactor. To run the
carbonylation and hydrogenolysis simultaneously in a single reactor, different combinations of
catalysts have been investigated, in particular CH
3
ONa/Cu and CH
3
ONa/Ni. Nickel-based systems are
very active and selective, but due to volatility and high toxicity of the Ni(CO)
4
that may arise during
the reaction, this process is considered difficult and hazardous to use in industrial settings. Copper
systems are thus preferred because they offer similar activities and selectivities, without the toxicity
problems associated with the nickel systems.
Another report states that continuous operation in a bubble reactor at 110
o
C and a pressure of only 5
atm is also possible. The problem with this process is the presence of CO
2
and water in the syn-gas,
reacting with sodium methoxide, thus deactivating the catalyst, and forming undesirable by-products.
To minimize this deactivation, CO
2
and water must therefore be removed from the gas feed. Catalysts
more tolerant to them (such as the recently reported KOCH
3
/Cu
2
Cr
2
O
5
systems) are also being
developed.
[31]
19
Although further improvements are still needed, the synthesis of methanol via methyl formate at
modest temperature and pressure could lead to an attractive alternative to current processes operating
at 200 300
o
C and 50 100 atm. At the same time, methyl formate can also be produced from the
hydrogenative conversion of CO
2
into formic acid and methanol as well as by the dimerization of
formaldehyde. The methyl formate to methanol route can also play a role in the secondary treatment
of the oxidative conversion of methane into methanol, without prior syn-gas formation (vide infra).
Recently developed indirect routes to methanol from: Catalytic hydrogenation of organic
carbonates, carbamates, and formates are of significant interest both conceptually and practically,
because these compounds can be produced from CO
2
and CO, and their mild hydrogenation can
provide alternative, mild approaches to the indirect hydrogenation of CO
2
and CO to methanol, an
important fuel and synthetic building block.
[32 33]
Although, it is believed that the lower electrophilicity of the carbonyl group, as a result of resonance
effects involving alkoxy or amido groups, makes hydrogenations of carbonic acid derivatives, such as
organic carbonates, carbamates and formates very difficult. Di-methyl carbonate [(MeO)
2
CO], a
stable and green solvent, is used in the reactions involving metal-catalyzed transformation.
Milstein and Co-workers
[33]
recently reported that (MeO)
2
CO can be selectively hydrogenated to
methanol, under rather mild conditions (T = 110145
o
C, P =10 60 atm., t
r
= 1 14 hrs). The success
of this process arises from the use of ruthenium (II) PNN pincer complexes A and B: derived from
pyridine-and bi-pyridine-based tridentate ligand, which are able to provide a hydride and a proton by
metal-ligand cooperation.
Also, the subsequent conversion of methyl formate to methanol, however, requires heterogeneous
hydrogenation in the gas-phase, which suffers from a lack of selectivity, and requires quite harsh
conditions. Interestingly, an unprecedented homogeneously-catalyzed hydrogenation of organic
formates to methanol has, recently, been accomplished. "A possible mechanism involves the metal-
ligand cooperation by aromatization-and de-aromatization of the hetero-atomic pincer core and
hydride transfer to the carbonyl group.
These atom-economical reactions proceed under neutral homogeneous conditions, at mild
temperatures, under mild hydrogen pressures, and can operate in the absence of solvent with no
waste-generation (green reaction). In spite of these remarkable achievements, more economical and
bio-relevant metal-based catalytic systems are required.
20
1.2.4
METHANOL PRODUCTION FROM SYN-GAS
Today, methanol is almost exclusively produced from syn-gas: a mixture of hydrogen, carbon
monoxide and some CO
2
, over a heterogeneous catalyst according to [Eq. (1.10) (1.12)]:
CO + 2H
2
CH
3
OH
H
0
298
= 21.7kcal mol
-1
(1.10)
CO
2
+ H
2
CO + H
2
O (in reverse)
H
0
298
= 9.8kcal mol
-1
(1.11)
CO
2
+ 3H
2
CH
3
OH + H
2
O
H
0
298
= 11.9kcal mol
-1
(1.12)
Since the first two reactions are exothermic (resulting in decrease of volume as the reaction proceeds),
control of the process temperature is important to avoid rapid deactivation of the catalyst.
[34]
Syn-gas for methanol production can be obtained by reforming or partial oxidation of any available
carbonaceous material such as coal, coke, natural gas, petroleum, heavy oils, and asphalt. Economic
considerations dictate the choice of the raw material. However, the long-term availability of raw
materials, energy consumption, and environmental aspects, also play important roles. The gas
treatment before methanol synthesis or addition of hydrogen is therefore needed to avoid the
formation of undesired by-products. Furthermore, natural gas contains fewer impurities, such as sulfur
and halogenated compounds, which can poison the needed catalysts. When present, these impurities:
mostly sulfur in the form of hydrogen sulfide (H
2
S), carbonyl sulfide (COS), or mercaptans, can
however, be removed relatively easily. Lower levels of impurities in the syn-gas allowed the use of
more active catalysts, operating under milder conditions.
Aside from the methanol synthesis step, the most crucial part of present methanol plants is the syn-
gas generation and purification system, which depend on the nature and purity of the feedstock used.
Although natural gas is generally the preferred feedstock due to the simplicity of obtaining an
adequate syn-gas with low levels of impurities, other routes are also used under given circumstances.
21
Fig. 1.2
A Scheme for Methanol Production via Natural Gas.
[35]
10 Reformer, 11 Reaction tube, 12 Combustion device, 13 Convection portion, 14 Chimney,
20
1
Fuel feeding passageway, 20
2
Raw gas-feeding passageway, 20
3
Water vapor-feeding
passageway, 20
4
Combustion exhaust gas-feeding passageway, 20
5,6,7,8,9,11,12,13,14,15,16,19,20,21,22
Passageways, 20
10
Exhaust passageway, 20
17
Purge gas passageway, 20
18
Branched purge gas
passageway, 30 CO
2
recovery apparatus, 31 Cooling column, 32 CO
2
Absorption tower, 33
Absorbing liquid regenerating tower, 34 Gas-liquid contacting member, 35a & 37aUpper gas-
liquid contacting member, 35b & 37b Lower gas-liquid contacting member, 36 Overflow portion,
38 Blower, 39;41;42 Pump, 40;43;44;52 Heat exchanger, 51;54 Compressor, 53 Heat recovery
device, 60 Methanol-synthesizing reactor apparatus, 61 Pre-heater, 62 Circulating passageway,
63 Methanol-synthesizing reactor, 71 Gas-liquid separator, 72 Cooling heat exchanger, 73 Gas-
circulating passageway, 74 Gas compressor, 80 Boiler, 90 Distillation apparatus.
1.2.4.1
SYNTHESIS-GAS PREPARATION TECHNIQUES
Syngas can be produced in a number of ways. The method of choice depends on the desired
characteristics of the syngas product. The principal routes of syngas production include natural gas,
petroleum oil and higher hydrocarbons, coal gasification, combined reforming, and heat reforming.
These syngas production routes have different operating costs and complexities. In general, there is no
22
overall best route, and project/site specifics will determine which syngas generation route is
preferred.
[22]
A measure of synthesis gas composition is given by the stoichiometric ratio (R)
mol. H
2
R =
(2
× mol. Co + 3 × mol. CO
2
)
(1.13)
Syngas which is hydrogen-rich has R values > 1.0. Hydrogen-lean mixtures have R values < 1.0. The
methanol synthesis plant performance characteristics will define the optimum R value for peak
methanol production. A brief review of the different syngas preparation techniques is presented below.
1.2.4.1.1
SYN-GAS VIA NATURAL GAS
1.2.4.1.1.1
STEAM REFORMING OF METHANE (SMR)
Natural gas is an appropriate feedstock for a synthesis plant containing Cu-Zn catalyst because it
commonly possesses low sulfur content. Simply stated, the natural gas feed is desulfurized, mixed
with steam, pre-heated at high temperatures (800 1000
o
C) and under pressure (20 30 atm.) over Ni
catalyst to form CO and H
2
,
[36]
which is then cooled.
A part of the CO formed reacts consequently with steam in the water-gas-shift-reaction (WGSR) to
yield more H
2
and also CO
2
. The gas obtained is thus a mixture of H
2
, CO and CO
2.
CH
4
+ H
2
O CO + 3H
2
H
0
298
= 49.1kcal mol
-1
(1.14)
CO + H
2
O CO
2
+ H
2
H
0
298
= 9.8kcal mol
-1
(1.15)
The ratio of the components depends on the reaction conditions: temperature, pressure and H
2
O/CH
4
ratio. The efficiency of syn-gas generation from methane, however, increases with increasing
temperature and decreasing pressure. With increasing temperatures, the WGS reaction becomes also
less dominant and the main products are CO and H
2
. Methane reforming can also be affected by the
thermal "coking" process involving the formation of carbon, which deposits as soot or coke on the
catalyst, internal parts of the reformer, and the downstream equipment, resulting in possible clogging.
Carbon may be formed by CH
4
decomposition or CO disproportionation (Boudouard reaction):
CH
4
C + 2H
2
H
0
298
= 18.1kcal mol
-1
(1.16)
2CO C + CO
2
H
0
298
= 40.8kcal mol
-1
(1.17)
In practice, the undesired carbonation, or soot formation, is largely prevented by the use of excess
steam and short residence times in the reactor. It can, however, be more problematic for the partial
oxidation process operating at higher temperature than methane reforming.
Since the steam reforming is highly endothermic high temperature, low pressure, and high steam-
to-carbon ratios enhance performance. Consequently, it is very energy intensive, expensive, and
23
represents the major expenditure in methanol production. Elimination of this step would provide a
significant economic advantage in methanol synthesis
1.2.4.1.1.2
CATALYTIC PARTIAL OXIDATION OF METHANE (CPO)
Partial oxidation is the reaction of methane with insufficient oxygen, which can be performed with or
without a catalyst.
[37]
This exothermic-reaction, operated at higher temperatures (T = 800 1500
o
C),
generally yields syn-gas with a H
2
/CO ratio = 2, which is ideal for methanol synthesis. The problem
with partial oxidation is that the products (CO and H
2
), can be further oxidized to form undesired
CO
2
and water in an highly exothermic reactions, raising safety concerns and leading to S values
typically below 2:
CH
4
+ ½O
2
CO + 2H
2
H
0
298
= 8.6kcal mol
-1
(1.18)
CO + ½O
2
CO
2
H
0
298
= 67.6kcal mol
-1
(1.19)
H
2
+ ½O
2
H
2
O
H
0
298
= 57.7kcal mol
-1
(1.20)
The production of excess energy in the form of heat is not desirable and wasteful if no immediate use
for it can be found for other processes.
1.2.4.1.1.3
AUTO-THERMAL REFORMING (ATR)
To produce syn-gas without either consuming or producing much excess heat, modern plants usually
combine exothermic partial oxidation with endothermic steam reforming to have an overall
thermodynamically-neutral reaction while obtaining syn-gas with a composition suited for methanol
synthesis (S
2). This process is called "auto-thermal reforming".
CH
4
+
3
/
2
O
2
CO + 2H
2
O
(1.21)
CH
4
+ ½O
2
CO + 2H
2
H
0
298
= 8.6 kcal mol
-1
(1.22)
CO + H
2
O CO
2
+ H
2
H
0
298
= 9.8 kcal mol
-1
(1.23)
Partial oxidation and steam reforming can be conducted simultaneously in the same reactor by
reacting methane with a mixture of steam and oxygen. Having only one reactor lowers the cost and
complexity of the system. However, because the two reactions are optimized for different temperature
and pressure conditions, they are generally conducted in two separate steps. After the steam reforming
step, the effluent from the reformer's outlet is fed to the partial oxidation reactor where all the residual
methane is consumed.
[38]
The oxygen required for the oxidation step means that an air separation plant
is needed, but to avoid the construction of such a unit, the use of air rather than oxygen is possible.
The produced syn-gas in this case will, however, contain a large amount of nitrogen and require
24
special processing before conversion into methanol. Thus, most modern methanol plants use pure
oxygen.
1.2.4.1.1.4
SYN-GAS FROM CO
2
REFORMING OF METHANE (CMR)
Syn-gas can also be produced by the reaction of carbon-dioxide with methane or natural gas,
generally termed CO
2
or "dry-reforming" because it does not involve any steam. With a reaction
enthalpy of (H = 59.1 kcal mol
-1
), this reaction is more endothermic than the steam reforming.
[39]
CH
4
+ CO
2
2CO + 2H
2
H
0
298
= 59.1kcal mol
-1
(1.24)
The reaction is carried out commercially at temperatures around 800 1000
o
C using catalysts based
on nickel (Ni/MgO, Ni/MgAl
2
O
4
). The syn-gas produced has a H
2
/CO ratio of 1, which is much lower
than the values of ~ 3 obtained with steam reforming.
While this lower ratio is a disadvantage for methanol synthesis, it makes a suitable feed gas for
other processes, especially iron ore reduction and FischerTropsch hydrocarbon synthesis.
For methanol production, hydrogen generated from other sources would have to be added to the
obtained syn-gas. Combination of steam and CO
2
reforming, which can also be performed using the
same catalysts to reach adequate syn-gas composition is also possible (bi-reforming, vide infra).
1.2.4.1.2
SYN-GAS FROM PETROLEUM OIL AND HIGHER HYDROCARBONS
Crude oil, heavy oil, tar and asphalt can all be transformed into syn-gas. This reaction involves partial
oxidation of heavy oils at temperatures of 1350 1600
o
C and pressures of up to 15MPa (150 atm).
However, this process is accompanied by incomplete combustion of heavy hydrocarbon feedstock
according to the following reactions [Eq. (1.25) (1.26)]:
C
n
H
m
+ nH
2
O n CO + (n +
m
/
2
) H
2
(1.25)
C
n
H
m
+
n
/
2
O
2
n CO + (
m
/
2
) H
2
(1.26)
A major advantage of this process is the usage of heavy feedstock not usable in other, vapor-only
processes. Disadvantages of the process include soot formation, pure oxygen requirement (instead of
air), high effluent levels of sulfur and sulfur derivatives (which can very rapidly poison the catalysts
used for steam reforming and for subsequent methanol synthesis). Considerably, more capital must be
invested in the purification steps, and catalysts that are more resistant to poisoning may be employed.
Several large refining companies (including Shell Oil Co. and Texaco Inc.) have achieved successful
commercialization of this type of partial oxidation process.
25
1.2.4.1.3
SYN-GAS FROM COAL
Syn-gas is produced from coal by gasification, a process combining partial oxidation and steam
treatment, according to the following reactions [Eq. (1.27) (1.30)]:
C + ½O
2
CO
H
0
298
= 29.4kcal mol
-1
(1.27)
C + H
2
O CO + H
2
H
0
298
= 31.3kcal mol
-1
(1.28)
CO + H
2
O CO
2
+ H
2
H
0
298
= 9.8kcal mol
-1
(1.29)
CO
2
+ C 2CO
H
0
298
= 40.8kcal mol
-1
(1.30)
Different coal gasification processes have been developed and commercialized over the years. The
selection of a particular design depends greatly on the characteristics of the coal used such as lignite,
sub-bituminous, hard coal and graphite have differing ash contents, moisture contents, levels of
impurities, caking behavior, reactivity, particle size distribution, and fixed carbon availability.
[6]
Owing to the low H/C ratio of coal, the obtained syn-gas is rich in carbon oxides (CO and CO
2
), but
deficient in hydrogen. Very low levels of some impurities (potassium, iron, etc.) can create a catalytic
environment, thus, modifying reaction rates. Often, effluent gas from a coal gasification unit requires
additional treatment before it can be used for methanol synthesis.
Impurities, especially sulfur and sulfur derivatives, must be removed to preclude poisoning of the
catalyst. Coal gasification yields a raw gas that is very carbon-rich (R < 1.0) and composition
adjustments are frequently necessary.
Large indigenous coal reserves were predicted to create a boom in the use of coal as a chemical
feedstock. However, the expense of coal conversion, variation in coal characteristics, and continued
availability of natural gas feedstock have prevented large-scale conversion to coal.
1.2.4.1.4
SYN-GAS VIA COMBINED REFORMING
Combined reforming (also known as "combination reforming" or "oxygen-enhanced reforming")
makes use of two reformers in series for syngas production. The primary reformer is operated
similarly to the natural gas reformer above, but the secondary (or auto-thermal) reformer is injected
with pure (99.5%) oxygen. Pure oxygen in the second reformer precludes the burden of compressing
large amounts of nitrogen. The oxygen also consumes excess hydrogen, making it possible to produce
a nearly stoichiometric syngas from the natural gas feedstock.
The advantage of this approach is brought about by shifting a portion of the reformer duty from the
primary reformer to the secondary reformer. The partial combustion that occurs in the secondary
reformer heats the process stream and allows reduction of the fired duty in the primary reformer.
26
In general, this approach is more costly than the steam reforming of natural gas approach, but is
justified in cases where energy costs are extremely high. Combined reforming does offer significant
environmental benefits over other approaches including reductions in CO
2
and NO
x
emissions.
[6]
Although steam reforming remains the main economical industrial process of syn-gas production,
processes combining partial oxidation with steam and carbon dioxide reforming are widely used to
reduce energy consumption, cut equipment cost, and obtain the optimal ratio between CO and H
2
.
1.2.4.1.5
SYN-GAS VIA HEAT-EXCHANGE REFORMING
In heat-exchange reforming, a heat-exchange reformer is operated in series with an auto-thermal
reformer. The central concept in this approach is that heat generated in the secondary or auto-thermal
reformer is used to heat the process gas reacting within the heat-exchange reformer. This approach is
simple, effective, and can provide syngas of nearly stoichiometric composition.
Heat-exchange reforming provides the following advantages: increased operational flexibility, high
reliability, reduced maintenance and energy costs, physically compact units, and reduced hazardous
emissions. Although both combined reforming and heat-exchange reforming are more efficient than
steam reforming, steam reforming still provides the most economical means of syngas generation for
methanol production.
1.3
METHANOL APPLICATION AND ECONOMY
Methanol serves three key markets: (i) Chemical, (ii) Transportation, and (iii) Power generation.
The dynamics of renewable methanol in each of these markets will depend on their respective
demand and prices. Market segments that offer an advantage for renewably-produced methanol such
as transportation and electricity generation, are more likely to see renewable methanol use, whereas,
market segments in which no such advantages are offered are expected to use the least expensive form
of methanol, likely produced from natural gas or coal.
1.3.1
METHANOL AS A CHEMICAL FEEDSTOCK
Methanol is also a starting material in the chemical and petrochemical industries for the production of
chloro-methanes, methylamines, methyl methacrylate, dimethyl terephthalate, and etc. These chemical
intermediates are then processed to manufacture many products of our daily life
[Fig. 1.3]
, such as: (i)
polymers, (ii) plywood, (iii) paints, (iv) pesticides, (v) plastics, (vi) resins, (vii) silicones, (viii)
adhesives, (ix) antifreeze in pipelines, (x) explosives, (xi) construction materials, (xii) magnetic film,
27
(xiii) windshield fluids, (xiv) as a lab solvent in high pressure liquid chromatography mass
spectroscopy (HPLC-MS) and UV/VIS spectroscopy, (xv) denaturant for ethanol and polyacrylamide
gel electrophoresis, and a potentially long list of other petrochemical products.
[40]
Fig. 1.3
Methanol Derived Chemical Products & Materials.
[40]
28
1.3.2
METHANOL AS A TRANSPORTATION FUEL
The knowledge gained with the in-house research and the demonstration fleets in the field has brought
the technology for methanol-fueled vehicles to a high level of reliability.
A recent study
[41]
by the Massachusetts Institute of Technology (MIT) compares liquid fuels into
which natural gas can be converted at room temperature, atmospheric pressure, thus recommends the
use of methanol (produced for a long period at large industrial scale) as a liquid natural gas carrier.
Liquid fuels have a better volumetric energy density than gaseous fuels. If spilled or leaked, methanol
is biodegradable and less ecologically damaging than conventional fuels.
[42]
They also are the most
compatible fuels with existing distribution systems and engines i.e. they require the least departure
from the technologies in place today for both the vehicle and the refueling infrastructure.
For the tank-to-wheels (TTW) portion of full fuel cycle, methanol as a transportation fuel can offer
two types of advantages over conventional fuels: (i) lower tailpipe emissions in combustion, and (ii)
higher efficiency vehicle technologies.
1.3.2.1
METHANOL AS FUEL IN INTERNAL COMBUSTION ENGINES (ICE)
Relative to gasoline, methanol has a higher octane rating, which also, makes it possesses an excellent
combustion characteristics that are ideal fuels for today's internal combustion engine (ICE)-driven
vehicles.
[43]
This higher octane rating allows the engine to run at a higher compression ratio (10 11
to 1 against 8 9 to 1 for gasoline engines) and thus, also more efficiently than a gasoline powered
engine. Efficiency is also increased by methanol's higher flame speed, which enables faster and more
complete fuel combustion in the cylinders.
Also, methanol has a latent heat of vaporization which is ~ 3.7 times higher than gasoline, so that
methanol can absorb a much larger amount of heat when passing from the liquid to gaseous state. This
helps to remove heat away from the engine so that it may be possible to use air-cooled radiators
instead of heavier, water-cooled systems. For similar performance to a gasoline-powered car, a
smaller, lighter engine block, reduced cooling requirements, better acceleration and mileage are to be
expected from methanol-optimized engines in the future.
[44, 49]
Methanol-specific engines, however,
provide even better fuel economy.
[45 46]
In addition, methanol vehicles have low overall emissions of
air pollutants such as hydrocarbons, NO
x
, SO
2
and particulates. Thereby, establishing an infrastructure
for methanol fuels would also greatly ease the introduction of advanced fuel cell-based vehicles using
methanol as a fuel either by on-board reforming to hydrogen, or directly with direct methanol fuel
cells, thus increasing its use will fundamentally change their need and production capabilities.
29
With diminishing oil and gas reserves, a new realization for the need of finding alternative solutions
is finally achieving a foothold, and the future of methanol (a high performance, a high efficiency fuel)
as a transportation fuel is entering a new period.
1.3.2.2
METHANOL AS FUEL IN COMPRESSION IGNITION ENGINES (DIESEL)
Particulate matter whether carcinogenic compounds or not, has been identified as a significant
health hazard, especially in large cities. Diesel fuel generally produces particles during combustion.
When combusted, methanol does not produce smoke, soot or particulates. This, and the fact that
methanol produces very low emissions of NO
x
because it burns at lower temperatures, makes
methanol attractive as a substitute for diesel fuel.
[45, 47]
Compression ignition engines (diesel) are quite different from gasoline engines. Instead of using
sparkplugs to ignite the fuel/air mixture in engine's cylinders, diesel motors rely on self-ignition fuel
properties to ignite under specific high-temperature and high-pressure conditions.
Methanol, also, has a significant vapor pressure compared to diesel fuel. The higher volatility allows
heavy-duty engines to start easily in the coldest weathers, thereby avoiding the white smoke that is
typical of cold starts with conventional diesel engines.
[45 46]
1.3.3
METHANOL AS AN ANTIKNOCK AGENT (GASOLINE ADDITIVE)
Methanol serves as an antiknock agent either in its pure or upgraded form to oxygenates
[47 48]
such
as MTBE (methyl tert-butyl ether), ETBE (ethyl tert-butyl ether), TAME (tert-amyl methyl ether). In
order to achieve an acceptable octane number, these oxygenates are added to gasoline to ensure
clean combustion, promote low tendency to vapor lock, and low engine deposit. Relative to other
oxygenates, MTBE, has many attractive properties including low heat of vaporization, low blending
Ried Vapour Pressure, and insensitivity to water.
Adoption of automobile fuel that are able to burn in neat form (100% pure) as a gasoline substitute,
or as a gasoline blend such as M85 would be highly beneficial as per environmental and energy
consideration, as well as necessitates a large increase in the World's methanol production capacity.
In a most recent study of tri-fuel blends containing varying levels of methanol, ethanol, and gasoline,
alcohol fuel blends show approximately 15 20 percent less carbon than gasoline.
[47, 50]
30
1.3.4
METHANOL FOR STATIC POWER AND HEAT GENERATION
Methanol is also an attractive fuel for static applications. It can be used directly as a fuel in gas
turbines to generate electric power.
[51]
Gas turbines use typically either natural gas or light petroleum
distillate fractions as fuels. Compared to these fuels, tests conducted by many institutions, beginning
in the 1970s, have shown that methanol can achieve higher power output and lower NO
x
emissions
due to lower flame temperatures. Since methanol does not contain sulfur, SO
2
emissions are also
eliminated. Operation using methanol offers the same flexibility as using natural gas and distillate
fuels, including the ability to start, stop, accelerate and decelerate rapidly, following the electric
power needs.
As an alternative to combustion, methanol can also be used in fuel cells, an advanced vehicle
technology where the fuel is used to generate electricity for motive power. DMFC are already being
used for producing electricity
[52 53]
in facilities sensitive to power outages such as airports, hospitals,
military complexes, and banks. For instance proton exchange membrane fuel cells are estimated to be
2.6 3.5 times more efficient than combustion engines
.
When fuel consumption is reduced, the
resulting tailpipe emissions are reduced as well. The wheels to tank (WTT), and tank to wheels (TTW)
GHG emissions of renewable methanol represent compelling benefits that can enable the
transportation sector to meet ambitious carbon reduction goals.
1.3.5
METHANOL USE FOR WASTEWATER DENITRIFICATION
Methanol is used by municipal and private wastewater treatment facilities to aid in nitrogen removal
from effluent streams.
[54]
As wastewater is collected in a treatment facility, it contains high level of
ammonia. Through a bacterial degradation process, this ammonia is converted into nitrate. If
discharged into the environment, the nutrient rich nitrate in sewage effluent can have a devastating
effect on water ecosystems consequently creating miles long algae blooms that sap oxygen and
sunlight from aquatic life. Advantages of methanol in the denitrification process are copious
[54]
i.e. it
contains no solids, no additional nutrients, has a neutral pH, contains 100% readily biodegradable
substrate, cost-effective (by revitalizing waterways tainted by the effects of nitrates).
1.3.6
METHANOL USE FOR BIODIESEL TRANS-ESTERIFICATION
In the process of making bio-diesel fuel, methanol is used as a key component in a process called
trans-esterification - to put it simply, methanol is used to convert the triglycerides in different types of
oils into usable bio-diesel fuel.
[55]
The trans-esterification process reacts methanol with the
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- Citar trabajo
- Ayodeji Ijagbuji (Autor), 2015, The Eco-Friendly and Promoting Influence of Nitric Acid–Steam Vapors on the Oxidation of C3–C4 Parrafins into Methanol, Múnich, GRIN Verlag, https://www.grin.com/document/288794
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