Chemical Kinetics of Ammonia Plant Reactions

Ammonia Plant
Chemical reaction
Kinetics
Praveen Singh

Introduction
CHEMICAL KINETICS
 Branch of chemistry
concerned with the rate of
chemical reactions and the
mechanisms by which
chemical reactions occurs
2

Chemical Reaction : Thermodynamics vs .Kinetics
• Thermodynamics says NOTHING about the rate of a reaction.
• Thermodynamics : Will a reaction occur ?
• Kinetics : If so, how fast ?
• Whether a reaction can proceed or not and to what extent a reaction can proceed is solely determined by the
reaction thermodynamics, which is governed by the values of DG & Keq, NOT by the presence of catalysts.
• In another word, the reaction thermodynamics provide the driving force for a rxn; the presence of catalysts changes
the way how driving force acts on that process.
3

Kinetic Vs. Thermodynamic
A reaction may have a large, negative DGrxn, but the rate may be so slow that there is no evidence of it occurring.
Example: Conversion of graphite to diamonds is a thermodynamic favor process (ΔG= -ve )
C(graphite) C(diamond)
Kinetics makes this reaction nearly impossible
(Requires a very high pressure and temperature over long time)
4

Kinetic Vs. Thermodynamic
• e.g CH4(g) + CO2(g) = 2CO(g) + 2H2(g) DG373=151 kJ/mol (100 °C)
DG973 =-16 kJ/mol (700 °C)
• At 100°C, DG373=151 kJ/mol > 0. There is no thermodynamic driving force, the reaction won’t proceed with or
without a catalyst
• At 700°C, DG373= -16 kJ/mol < 0. The thermodynamic driving force is there. However, simply putting CH4 and CO2
together in a reactor does not mean they will react. Without a proper catalyst heating the mixture in reactor
results no conversion of CH4 and CO2 at all. When Ni/Al2O3 is present in the reactor at the same temperature,
equilibrium conversion can be achieved.
5

WHAT IS EQUILIBRIUM OF A REACTION?
EQUILIBRIUM OF A REACTION IS
For a reversible chemical reaction , equilibrium means a state where the forward
and backward reaction rates are equal and there is no change in composition of
the system with time .
Thermodynamically , equilibrium is a state where the total Gibb’s free energy (Gt)
of all the species of a system is at the minimum under given temperature(T) and
pressure(P) conditions .
( dGt)T,P = 0 G = H - T*S =
Energy available for useful work ( H = Enthalpy , S = Entropy )

Order and
Molecularity of
a reaction
Molecularity of an elementary reaction is
defined as number of molecules involved in
the reaction and it will always be an
integer.
Order refers to empirically found rate
expression and it can have a fraction value
need not be an integer.
7

Le Chatelier’s
Principle
If a stress is applied to a system in
dynamic equilibrium, the system
changes to relieve the stress.
System stresses:
• Concentration of reactants or
products
• Temperature
• Pressure
8

Factors That Affect Reaction Rates
• Concentration of Reactants
• As the concentration of reactants
increases, so does the likelihood that
reactant molecules will collide.
• Temperature
• At higher temperatures, reactant
molecules have more kinetic energy,
move faster, and collide more often and
with greater energy.
• Catalysts
• Speed reaction by changing mechanism.
9

Effect of Concentration of Reactants
 Adding reactant shifts the reaction toward the
products.
 Stress: Increasing reactants
 Relief: Decreasing reactants
 Shift: to the right (products)
H2O (l) + CO2 (g) H2CO3 (aq)
10

Effect of Concentration of Products
 Adding products shifts the reaction toward the
reactants.
 Stress: Increasing products
 Relief: Decreasing products
 Shift: to the left (reactants)
H2O (l) + CO2 (g) H2CO3 (aq)
11

Effect of Temperature
 Increasing the temperature causes the
equilibrium to shift in the direction that absorbs
heat.
 Stress: Increase in Temp
 Relief: Decrease in Temp
 Shift: Towards the left
SO2 (g) + O2 (g) 2SO3 (g) + heat
12

Effect of Pressure
 Affects gases only.
 For unequal number of moles of reactants and
products, if pressure is increased, the
equilibrium will shift to reduce the number of
particles.
 For equal number of moles of reactants and
products, no shift occurs.
2NO2 (g) N2O4 (g)
13

Methods for measuring reaction rate
 Since the reaction rate is the change in the
amount of a product or a reactant per unit time,
any property that is related to amount of
product or reactant present can be used to
measure the rate of reaction. Some properties
that can be used to measure the reaction rate
are given in the following table.
14

Classification
of Chemical
Reaction
There are following type of chemical
reaction-
a) Homogeneous Reaction
b) Heterogeneous Reaction
c) Single reaction
d) Multiple reaction
e) Elementary reaction
f) Non- elementary reaction
16

The Collision Model
• In a chemical reaction, bonds are broken and
new bonds are formed.
• Molecules can only react if they collide with
each other.
• Furthermore, molecules must collide with the
correct orientation and with enough energy to
cause bond breakage and formation.
17

Activation Energy
 There is a minimum amount of energy required
for reaction: the activation energy, Ea.
 Just as a ball cannot get over a hill if it does not
roll up the hill with enough energy, a reaction
cannot occur unless the molecules possess
sufficient energy to get over the activation
energy barrier.
18

Reaction Coordinate Diagrams
• It shows the energy of the reactants and
products (and, therefore, DE).
• The high point on the diagram is the transition
state.
• The species present at the transition state is
called the activated complex.
• The energy gap between the reactants and the
activated complex is the activation energy
barrier.
19

Arrhenius Equation
Svante Arrhenius developed a mathematical
relationship between k and Ea:
where A is the frequency factor, a number that
represents the likelihood that collisions would
occur with the proper orientation for reaction.
20

Catalysts
• Catalysts increase the rate of a reaction by
decreasing the activation energy of the
reaction.
• Catalysts change the mechanism by which the
process occurs.
21

• They alter the potential energy of the
intermediate complex, thus making their
formation easier / slower .
• Activity of catalysts is very sensitive to
temperature .
• Catalysts are selective .
• They don’t affect the equilibrium
composition , only alter the time to reach
equilibrium .
22

General
requirements
for a good
catalyst
 Activity - being able to promote the rate of desired
reactions
 Selective - being to promote only the rate of desired
reaction and also retard the undesired reactions
Note: The selectivity is sometime considered to be more
important than the activity and sometime it is more
difficult to achieve (e.g. selective oxidation of NO to NO2
in the presence of SO2)
 Stability - a good catalyst should resist to
deactivation, caused by
 the presence of impurities in feed (e.g. lead in
petrol poison TWC)
 thermal deterioration, volatility and hydrolysis of
active components
 attrition due to mechanical movement or pressure
shock
 Surface Area-A solid catalyst should have reasonably
large surface area needed for reaction (active sites). This
is usually achieved by making the solid into a porous
structure.
23

Mechanism of Heterogeneous Catalytic
Reaction
• The long journey for reactant molecules to
j. travel within gas phase
k. cross gas-liquid phase boundary
l. travel within liquid phase/stagnant layer
m. cross liquid-solid phase boundary
n. reach outer surface of solid
o. diffuse within pore
p. arrive at reaction / active site
q. be adsorbed on the site and activated
r. react with other reactant molecules, either being adsorbed on the
same/neighbour sites or approaching from surface above
• Product molecules must follow the same track in the reverse direction to
return to gas phase
• Heat transfer follows similar track
24
j
r
gas phase
poreporous
solid
stagnant layer
k
l
mn
o
p q
gas phase
reactant molecule

Catalyst
Preparation
Unsupported Catalyst
Usually very active catalyst that do not require high
surface area
e.g., Iron catalyst for ammonia production (Haber
process)
Supported Catalyst
requires a high surface area support to disperse the
primary catalyst the support may also act as a co-catalyst
(bi-functional) or secondary catalyst for the reaction
(promoter)
E.g., Highly dispersed Nickel clusters on MgAl2O4
25

Catalyst carrier
Physical material that is neutral, such as alumina, activated carbon or
silica, that supports a catalyst by increasing its surface area.
Requirements:
• Should with stand to high temperature
• Should with stand to high steam partial pressure.
• Should provide more surface area.
• Should be inert with the catalyst metal and tube metal
Classification of carrier
Cement type:- These are having
i. High crushing strength
ii. Volatile impurities, Like silica.
iii. Significant interaction with nickel.
iv. Eg.:- Calcium aluminate silicate, Calcium aluminate.
Ceramic type:
i. Good crushing strength.
ii. Almost inert with catalyst
iii. High temperature resistance.
iv. Eg:- Alpha-Alumina, Magnesium aluminate.

Catalyst
Disintegration
The various types of catalyst disintegration phenomenon are described below.
1) Fouling:- Formation of carbonaceous deposits on the catalyst, blocking pores
and sometimes masking complete surface known as fouling.
2) Aeging:- This is least understood type of deactivation of catalyst and is due to
using the catalyst over a long period of time.
3) Milling:- The phenomenon of milling is the breakage of the catalyst into fine
particles or powder and crusting as a consequence of rapid swirling of catalyst
due to impingement of high velocity gas on the catalyst surface.
4) Retardation:- Deactivation of catalyst due to coverage of active sites either
by reactants (Reactant inhibition) or by products (Product inhibition)
5) Sublimation:- Sublimation of catalyst agents deposited on inert support due
to any hot spot developed in the reactor.
6) Poisoning:- Some compounds occupy active sites permanently which
decreases total available active sites. These compounds are called poisons
and phenomenon is called poisoning.
Example: Sulfur, arsenic, chlorine, silica, oxygen are some of the most common
catalytic poisons.
Catalytic poisons act in various ways. They occupy active sites on catalytic
surface, distort the crystalline structure, thereby reducing strength of the
catalyst. They also block the catalyst pores, thereby reducing the surface area
available.
7) Sintering:- Melting of metal crystals at corners and recrystallization which
results in loss of surface area due to change in metal ensembles available.

CATALYST
PERFORMANCE
MONITORING
TECHNIQUES
The below are the some of most common practices used in chemical industry
for catalyst performance monitoring on regular basis by a plant operator.
1. Monitoring temperature profile across bed – The differential temperature
across catalyst bed must be constant for the same inlet gas composition
2. Pressure drop – A less pressure drop indicates good strength of catalyst.
3. Outlet and inlet gas composition – This analysis monitoring is very
important for predicting catalyst performance and also future life expectancy.
4. Approach to equlibrium – An approach to equilibrium is the difference
between the reactor outlet temperature and the average bed temperature of
the reactor?. A less approach to equilibrium indicates good performance of
the catalyst.
5. Z-90 and Z-70 curves – These curves represents that 90% and 70% of the
reaction completed at what depth of the catalyst bed respectively. By
following these curves one can predict the remaining catalyst life expectancy.
(These curves especially used for pre reformer in ammonia plant).

Equilibrium &
Reaction
Kinetics
involved in
Production of
Ammonia

Major reactions leading to synthesis of
Ammonia
• Reforming of Hydrocarbons
• Shift Conversion
• Synthesis of Ammonia

Other reactions
• Sweetening of Hydrocarbons
• Absorption / Stripping of CO2
• Methanation of CO / CO2 in synthesis gas

Reforming of hydrocarbons
• A mixture of hydrocarbons reacts with steam , in presence of a catalyst to produce a
mixture of CO , CO2 , H2 , CH4 .
• The purpose is to liberate hydrogen contained in hydrocarbon and steam.
• It is carried out in stages
*Adiabatic pre-reforming
*Primary reforming
*Secondary reforming

Adiabatic pre-reforming
Reforming of higher hydrocarbons (C2+) is carried out at a lower temperature ( 460 - 490C ) ,
where overall conversion (to CO,CO2, H2 ) is low , but higher hydrocarbons are converted into CH4.
Advantages :
* Prevention of higher hydrocarbons being exposed to a higher temperature in primary reformer
reduces the possibility of carbon deposition .
* Requirement of steam/carbon ratio is reduced .
* Pre-reformer catalyst chemisorbs S-compounds , the lower operating temperature being favorable
for the equilibrium , thus acts as a sulfur-guard for primary reformer catalyst .
* Reforming capacity is enhanced .

Primary and secondary reforming
• Primary reforming : A higher conversion CH4 , into CO , H2 , is achieved operating at
a higher temperature ( around 800°C at reformer outlet ) .
• Secondary reforming : Partial combustion of process gas leaving primary reformer is
achieved by introducing air , which leads to increase in temperature (1100 - 1200°C ) .
Higher temperature favors reforming of CH4 , and almost complete conversion of
hydrocarbons is achieved . Also , air introduces N2 ( required for synthesis of NH3 )
into the process gas .

Reactions involved in reforming
1) Cn Hm + n H2 O n CO + ( n + m/2 ) H2 - heat ΔH°298 > 0
2) CH4 + H2O CO + 3 H2 - heat ΔH°298 = 206.2 k J / mol
3) CO + H2O CO2 + H2 + heat Δ H°298 = - 41.2 k J / mol
Δ H° for (2) > Δ H° for (3)
* In pre-reforming , overall reaction is endothermic in NG case , whereas it is exothermic /
thermoneutral in naphtha case.
* In primary reforming reaction (2) predominates

Partial combustion reactions takes place in secondary reformer , followed by reaction (2) .
4) 2 H2 + O2 2 H2O + heat ΔH°298 = - 242.175 k J / mol
5) CH4 + 2 O2 CO2 + 2 H2O + heat ΔH°298 = - 561.3 k J / mol
Side reactions :
6) 2 CO C + CO2 + heat ( Boudard’s reaction ) ΔH°298 = - 172.4 k J / mol
7) CH4 C + H2 - heat ΔH°298 = 74.9 k J / mol
Reactions (6) & (7) are undesirable as they lead to carbon deposition.

• Minor reactions :
8) CH4 + CO2 2 CO + 2 H2 - heat
D H298 = 247.27 k J / mol
9) 2 CH4 C2 H4 + 2 H2 - heat
D H298 = 202.09 k J / mol
10) N2 + 3 H2 2 NH3 + heat
D H298 = - 45.96 k J / mol

Reaction equilibria for reforming
• Steam reforming of CH4 and other hydrocarbons , [reactions (1) & (2) ] , being
endothermic reactions are favored by higher temperature , Similarly, reactions (3) , (6)
[exothermic] are favored by lower temperature and reaction(7) [endothermic] is
favored by higher temperature .

• Equilibrium constant ( KP ) :
For reaction(2) , KP = ( pH2
3 * p CO ) / ( p CH4 * p H2O ) ( Favored by low pressure )
For reaction(3) , KP = ( p CO2 * p H2 ) / ( p CO * p H2O ) ( Not affected by pressure )
For reaction(6) , KP = ( p CO2 ) / ( p CO
2 ) ( Favored by higher pressure )
For reaction(7) , KP = ( p H2
2) / ( p CH4) ( Favored by lower pressure )

Approach to Equilibrium in reforming
• Equilibrium approach : It refers to the closeness of reactor exit composition to the
equilibrium composition at the catalyst exit temperature .
• ΔT( approach ) = T( K R) - T(catalyst exit)
K R = Reaction coefficient , calculated from actual composition put in equilibrium
constant expression .
ΔT(approach) is positive for exothermic reactions
negative for endothermic reactions

Mechanism / kinetics of reforming
• Ni acts as the active component of the reforming catalysts , while Mg O , Al2 O3 etc act as carrier .
• MgO reacts with steam at lower temperature (620 K / 347°C) reducing mechanical strength of catalyst .
Mg O + H2O Mg (OH)2 + heat
• Mechanism :
CH4 + A* CH2 - A* + H2 (11)
(Active site) (Active complex)
Formation of the active complex will be much easier for higher hydrocarbons .
CH2 - A* + H2O CO - A* + 2 H2 (12)
CO - A* CO + A* (13)
or , CH2 - A* C - A* + H2 (14)
C - A* + H2O CO + H2 + A* (15)

• CO + A* CO - A* (16)
CO - A* + H2O CO2 + H2 + A* (17)
or , H2O + A* O - A* + H2 (18)
O - A* + CO CO2 + A* (19)
Insufficient steam , may lead to ,
C - A* C + A* (20)
CO - A* + CO CO2 + C + A* (21)

• For reaction (2) ,
* Step (11) , is the rate determining step .
* The reaction is first order with respect to CH4
* An empirical rate expression ,
- r = [ 1.1* 109 exp(-15.6*103/T) /(1+a* p H2O/ pH2 + b* p H2O ) ] * p CH4
where, r is in mol m-3 h-1
a ,b constants depending on temperature .
Partial pressures are in M Pa .

Steam/ Carbon
ratio
A higher ratio gives ,
* Higher conversion of hydrocarbons .
* Less possibility of carbon deposition
.
A lower ratio gives ,
* Lower energy consumption.
* Lower operating temperature for
Shift Conversion .
The ratio has little effect on rate of
reactions .

Shift
Conversion
• Conversion of CO in process gas exit secondary
reformer , into CO2 is achieved by operating at
lower temperature , since the process is
exothermic .
CO + H2O CO2 + H2 + heat
• Limitations of operating at lower temperature :
* Lower reaction rate .
* Dew point of the process gas .
• The conversion of CO is achieved under various
combinations of operating temperatures.
* H . T . , M . T . , & L . T. Conversion
* H . T . & L . T . Conversion
* M . T . & L . T . Conversion

Shift
Conversion
• Operation at lower temperature is achieved by
* Lower Steam / Carbon ratio .
* Higher catalyst activity .
• M . T . S . / L . T . S are operated at 220 - 300C and below
200C respectively .
• Dew point of process gas at ,
M. T. S inlet = 183C (approx.) L. T. S inlet =
173C(approx..)
• Cu based catalyst for M . T . S. / L . T . S . Conversion :
* Cu is the active component .
* Al2O3 , ZnO act as carrier .
* ZnO also act as sulfur guard .
* High activity at lower temperature .

Reaction equilibria for Shift Reaction

Ammonia Synthesis
• Synthesis of ammonia takes place according to the reaction
N2 + 3H2 2NH3 + heat ΔH°298 = - 45.96 k J/ mol
• The exothermic reaction leads to depletion in the number of moles . So the favorable
conditions are * Higher pressure * Lower temperature
• Operation at a very low temperature is not convenient due to low rate of reaction .Higher
yield is achieved by:
• Choosing optimum operating temperature , pressure , circulation rate (space velocity) .
• Efficient product separation , heat recovery .
• High catalyst activity .

Reaction equilibria for Ammonia Synthesis
• The equilibrium constant is given as ,
• Kp = ( p NH3 ) / { ( p N2 )1/2 *( pH2 )3/2}
log Kp= 9591/(4.571T) - (4.98/1.985)logT-0.00046T/4.571 +(0.85*10-6 T2/4.571) + 2.1
where, T is Absolute temperature (K)
and, Partial pressures are in atmospheric unit .
• The heat of reaction is obtained as
DH = - 9500 - 4.96 T - 0.000575 T2 + 0.0000017 T3 from Haber’s data .DH is in Cal / mol.
As T increases Kp decreases , while - DH increases .

Reaction equilibria for Ammonia Synthesis

Kinetics of
Ammonia
Synthesis
• Synthesis of ammonia is carried out in
presence of Iron catalyst promoted by small
amount of oxides such as K2O , Al2 O3 .
• Free Fe on the surface layer of the iron oxide
base act as the active catalytic component . It is
formed by reduction of iron oxide with H2 .
Fe3O4 + 4 H2 3 Fe + 4 H2 - 34 Kcal
• Mechanism of catalytic synthesis of ammonia
is a complicated one , and not fully understood
. It seems a number of slow steps are involved .
The rate equation is affected also by
* Heat transfer rate
* Mass transfer (diffusion) rate

Kinetics of
Ammonia
Synthesis
• Mechanism : A surface covered by N atoms is
created by interaction of N2 molecules with
active catalyst component .
* N2 molecules are adsorbed on this layer of
atomic N . This step seems to be the main rate
determining step .
N* + N2 N*-- N2 ( adsorbed)
N2 ( adsorbed ) + 3H2 2NH3
(Intermediate steps)

Kinetics of
Ammonia
Synthesis
• The rate equation is given as ,
r = k1 pN2 ( p H2
3/ p NH3
2 )a - k2( p NH3
2/ p H2
3 )b
( forward reaction) ( backward
reaction)
k1 = rate constant for forward reaction
k2 = rate constant for backward reaction
a , b are empirical constants , where a + b = 1
These constants are dependent on ,
* Temperature
* Reactor volume
* Catalyst weight and surface area

Kinetics of
Ammonia
Synthesis
• Poisoning of synthesis catalyst : Oxygen and its
compounds ( e.g. CO , CO2 , H2O ) act as poison
to the ammonia synthesis catalyst .
CO and O2 are of special significance as the
others are separated out in the ammonia
separation step of the synthesis loop
operation .
• Mechanism of poisoning : These compounds
react with H2 to give H2O .
O2 + 2H2 = 2H2O
CO + 3H2 = CH4 + H2O
CO2 + 4H2 = CH4 + 2H2O
The H2O vapor takes part in the ammonia
synthesis mechanism leading to blocking of
active catalyst site by atomic oxygen.
N2(adsorbed) + H2O + 2H2 = 2NH3 + O(adsorbed)

Kinetics of
Ammonia
Synthesis
• Effect of catalyst poisoning of reaction rate
:
Presence of water vapor in the ammonia
synthesis system affect the reaction
mechanism in such a way that the reaction
is slowed down
r = k1 p N2{ p H2
2. 5/( p NH3 * p H2O )}- k2{p NH3 /(p
H2
0. 5*p H2O)
An increase in partial pressure of water
vapor leads to a slower rate .
• Sulfur , phosphorus are also severe poisons
to the synthesis catalyst . Carryover of oil
with the synthesis gas will also poison the
catalyst by clogging the micro-pores .

Effect of Space
Velocity
• Space velocity is a term related to reactor
design .
Space velocity = No. of reactor volumes( measured
under specified conditions ) of reacting fluid
processed per unit time ( hr -1 )
Space time = ( Space velocity )-1
Volume of fluids may be calculated under reactor
inlet conditions or normal conditions . Space
velocity is directly related to the circulation rate
.
• A high space velocity results in low conversion
through the reactor . But the decrease in
conversion is not proportional to the increase in
space velocit.

Effect of gas
composition
Effect of gas composition on
Equilibrium / Reaction rate can be
discussed under following points .
• Concentration of ammonia at
reactor inlet .
• Hydrogen / Nitrogen ( molal ) ratio .
• Concentration of inert gases .

Effect of gas
composition
(Ammonia
content at
reactor inlet )
• Thermodynamically , a higher concentration of
ammonia at the reactor inlet will result in
lower yield of ammonia due to equilibrium
constraints . So a lower concentration of
ammonia at the reactor inlet gives a higher
yield .
• As is evident from the rate equation , as
concentration of ammonia goes on increasing ,
rate of formation of ammonia goes on
decreasing . A higher concentration of
ammonia at the reactor inlet will give a low
initial rate ,which will further decrease. This
will also reduce the rate of temperature
increase due to heat of reaction , which will
again adversely affect the rate of reaction.
The final effect is a low yield . So a low
ammonia content at the inlet is favorable for
higher yield of ammonia .

Effect of gas
composition
( Effect of
H2/N2 Ratio )
• Maximum rate of formation of ammonia is
achieved when H2 / N2 molar ratio is
maintained constant at the stoichiometric
ratio ( 3 / 1 ) of the reactants .
• The rate of reaction starts falling rapidly
beyond the ratio 3.5/1 and 2.5/1 as
observed from experiment .
• To maintain the ratio at the reactor inlet
around 3/1 the ratio in the make up gas to
system is to be kept slightly lower to account
for difference in solubility of H2 and N2 in
ammonia .

Effect of gas
composition
( Content of
inert gases )
• In the ammonia synthesis gas , some gases are
present which do not take part in the
synthesis reaction , namely methane ( CH4 ) and
argon ( Ar ) . These are called inert gases .
• Under constant pressure conditions , larger
amount of inert will reduce the effective
pressure ( the partial pressure of reactants ) ,
adversely affecting the equilibrium yield , as
well as the reaction rate .
• If pressure is allowed to increase with
increasing inerts , this will result in higher
compression load .

Methanation
• The purpose of methanation is to convert traces
of carbon oxides present in ammonia synthesis
gas to methane , as such oxides are poisonous to
the synthesis reaction catalyst .
• The reactions are just the reverse of methane
reforming .
CO + 3 H2 CH4 + H2O + heat
ΔH°298 = -206.2 k J / mol
CO2 + 4 H2 CH4 + 2 H2O + heat
ΔH°298 = -165 k J / mol
The reactions are carried out on Ni catalyst .

Methanation
• The rate of the reactions are of
first order with respect to CO / CO2
, which are the limiting reactants
• H2 , present in large excess does not
affect the rate expression .
• The rate constant k is greater in
the reaction involving CO , under
same operating conditions .
Approximately , k CO = 2 k CO2

*
Hydrodesulfurization
*
• Purpose : Hydrocarbon feedstock ( NG or naphtha ) is treated
with H2 to remove sulfur content , which is poisonous to various
catalysts .
• Sulfur is removed in the form of H2S which is absorbed by a bed
of ZnO .
• RSH + H2 = RH + H2S (R may be any akyl group
RSR + H2 = RR + H2S like -CH3 )
RSSR + 2H2 = RR + 2H2S
COS + H2 = CO + H2S
H2S + ZnO = ZnS + H2O
COS + ZnO = ZnS + CO2

• These reactions are normally slightly
exothermic .
• The catalyst used for these reactions are Co-
Mo or Ni-Mo type .
• The operating temperature range is 380 -400°C
. Above 400°C , polymerization and deposition
may occur .
• The activity increases , when in sulfided state .
The active component is Mo-S2* .
• The sulfur in catalyst remains in equilibrium
with the sulfur in the gas phase .
• Unsulfided catalyst , if exposed to a few % of
CO2 , will lead to methanation , followed by
rapid rise in temperature . This requires the
catalyst to be exposed to N.G. ( carrying CO2) to
presulfided .

• Desulfurisation is possible even without
hydrogenation.
RH-S-HR = R=R + H2S - heat
• In the absence of sufficient H2, the unsaturated
compound may polymerize rapidly at the
operating temperature .
n R=R ( -R-R- )n ( without H2 )
R=R + H2 RH-RH ( with H2 ) + heat
( Here , R denotes an alkylene group like - CH2- )

Chemical Kinetics of Ammonia Plant Reactions

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