6.

1 REACTION EQUILIBRIUM
The key to understanding reaction equilibrium is the Gibbs free energy, or
free energy, A negative value of G implies a spontaneous reaction of
Before considering reaction conditions, some basic reactants to products. A positive value of G implies the reverse reaction is
(For Real system)
spontaneous.
principles of chemical equilibrium need to be
reviewed.
(T constant)
For reversible reactions:
a. For a given mixture of reactants at a given temperature and
(For ideal Gas)
pressure, there is a maximum conversion (the equilibrium
conversion) that cannot be exceeded and is independent of the
reactor design.
b. The equilibrium conversion can be changed by appropriate
(From P1 to P2)
changes to the concentrations of reactants, temperature and
pressure.
(Go at standart
pressure)

(for real
system)

1. Homogeneous gaseous reactions 2. Homogeneous liquid reactions

Because
For homogeneous gaseous reactions, the standard state fugacities can
be considered to be unity, that is, fio=1.

For ideal solution i =1

For an ideal gas φi = 1, thus:

(At equilibrium G=0)

Example 6.1 A stoichiometric mixture of nitrogen and hydrogen is to be reacted at

1 bar:

Assuming ideal gas behavior (R = 8.3145 kJ·K−1·kmol−1), calculate:

a. equilibrium constant
b. equilibrium conversion of hydrogen
c. composition of the reaction products at equilibrium at 300 K. Standard free
energy of formation data are given in Table 6.1
 6.2 REACTOR TEMPERATURE

because
(P Constant)

If H independence to T
From

at standard conditions for finite changes in Go and Ho

For more accurate H dependence to T
From definition

Example 6.2 Following Example 6.1:

a. Calculate ln(Ka) at 300 K, 400 K, 500 K, 600 K and 700 K at 1 bar and test the
validity of Equation 6.39. Standard free energy of formation and enthalpy of
formation data for NH3are given in Table 6.4. Free energy of formation data for
H2 and N2 is zero.

b. Calculate the values of ln(Ka) from Equation 6.39 and Equation 6.48 from standard
data at 298.15 K and compare with values calculated from Table 6.4. Heat capacity
coefficients are given in Table 6.5

c. Determine the effect of temperature on equilibrium conversion of hydrogen using

the data in Table 6.4.

where Again assume ideal gas behavior and R = 8.3145 kJ·K−1·kmol−1.

Because
Now consider the effect of temperature on the rate of reaction. 1. Single reactions. (b) Exothermic reactions. For single exothermic irreversible reactions, the
(a) Endothermic reactions. If an endothermic reaction is reversible, then Le temperature should be set as high as possible, consistent with materials of
Chˆatelier’s Principle dictates that operation at a high temperature increases the construction, catalyst life and safety, in order to minimize reactor volume.
maximum conversion. Also, operation at high temperature increases the rate of
reaction, allowing reduction of reactor volume. Thus, for endothermic reactions, For reversible exothermic reactions, the situation is more complex. Figure
the temperature should be set as high as possible, consistent with safety, 6.5a shows the behavior of an exothermic reaction as a plot of equilibrium
At difference temperature, and assume E constant
materials-of construction limitations and catalyst life conversion against temperature.

Generally, the higher the rate of reaction, the smaller is the reactor volume.
Practical upper limits are set by safety considerations, materials-of-construction
limitations, maximum operating temperature for the catalyst or catalyst life.
Whether the reaction system involves single or multiple reactions, and whether
the reactions are reversible, also affects the choice of reactor temperature.

Figure 6.4 Equilibrium behavior with change in temperature for endothermic reactions. Figure 6.5 Equilibrium behavior with change in temperature for exothermic reactions.

2. Multiple reactions. The arguments presented for minimizing reactor volume Example 6.3
for single reactions can be used for the primary reaction when dealing with Example 5.4 developed a kinetic model for the manufacture of benzyl acetate from
multiple reactions. However, the goal at this stage of the design, when benzyl chloride and sodium acetate in a solution of xylene in the presence of
dealing with multiple reactions, is to maximize selectivity or reactor yield triethylamine as catalyst, according to:
rather than to minimize volume, for a given conversion.

The reaction rate constants k1 and k2 both

increase with increasing temperature.
Example 5.4 developed a kinetic model for an equimolar feed at a temperature of As in Example 5.4, one way this can be carried out is by
102◦C, such that: setting up a function for R2 in the spreadsheet, and then
• If k1 increases faster than k2, operate at
using the spreadsheet solver to minimize R2 by manipulating
high temperature (but beware of safety, the value of kA (see Section 3.9). The results are given in
catalyst life and materials of Table 6.8.
construction constraints).
• If k2 increases faster than k1, operate at
low temperature (but beware of capital
cost, since low temperature, although Further experimental data are available at 117◦C. The measured mole per cent
increasing selectivity, also increases benzyl chloride versus time in hours at 117◦C are given in Table 6.7. Again, assume
reactor size). the volume of the reactor to be constant.

Determine the activation energy for the reaction.

6.3 REACTOR PRESSURE

- For reversible reactions, pressure can have a significant effect on the

equilibrium conversion.

- Even though the equilibrium constant is only a function of temperature

and not a function of pressure, equilibrium conversion can still be
influenced through changing the activities (fugacities) of the reactants
and products.
 7.1 TEMPERATURE CONTROL

- In the first instance, adiabatic operation of the reactor should be

considered since this leads to the simplest and cheapest reactor design.

By contrast with ideal models, practical reactors must - If adiabatic operation produces an unacceptable rise in temperature
for exothermic reactions or an unacceptable fall in temperature for
consider many factors other than variations in endothermic reactions, this can be dealt with in a number of ways:
temperature, concentration and residence time. a. Cold shot and hot shot.
b. Indirect heat transfer with the reactor.
c. Heat carrier.
d. Catalyst profiles.

7.2 CATALYST DEGRADATION 7.3 GAS–LIQUID AND LIQUID–LIQUID REACTORS

1. Gas–liquid reactors.
The performance of most catalysts deteriorates with time. The rate at which (In the gas film)
the deterioration takes place is not only an important factor in the choice of
catalyst and reactor conditions but also the reactor configuration. There are many reactions involving more than one reactant, where the reactants are
Loss of catalyst performance can occur in a number of ways: fed in different phases as gas–liquid or liquid–liquid mixtures.
a. Physical loss.
b. Surface deposits. This might be inevitable because the feed material is inherently in different phases at
c. Sintering (Sintering is a molecular rearrangement that occurs below the the inlet conditions.
melting point of the material and causes a reduction in the effective surface Alternatively, it might be desirable to create two-phase behavior in order to remove an
area of the catalyst.) unwanted component from one of the phases or to improve the selectivity. If the
d. Poisoning. reaction is two-phase, then it is necessary that the phases be intimately mixed so that
e. Chemical change. mass transfer of the reactants between phases can take place effectively. (In the liq film)

The overall rate of reaction must take account of the mass transfer resistance in order
to bring the reactants together as well as the resistance of the chemical reactions. The
three aspects of mixing, mass transfer and reaction can present widely differing Figure 7.2 The gas–liquid interface.
difficulties, depending on the problem.
If equilibrium

If steady state is assumed (NG,i = NL,i = Ni),

If kG,i is large relative to kL,i/Hi , the mass

transfer is liquid-film controlled.