OPEN CHEMICAL EQUILIBRIA AND COEFFICIENTS
OF THERMODYNAMIC ACTIVITY
B. Zilbergleyt,
Independent Scholar
E-mail: livent1@msn.com
ABSTRACT.
The lecture presents new approach to equilibrium in open chemical systems
suggesting linear dependence of the reaction shift from equilibrium upon
external thermodynamic force. Basic equation of the theory contains
traditional term with thermodynamic activities (logarithmic term in case of
constant equation) and a non-traditional, parabolic term resulting from
external impact due to system openness. At isolated equilibrium the nontraditional term equals to zero turning the whole equation to the traditional
form. The parabolic member coincides with the excessive thermodynamic
function and may be expressed in terms of system reaction to the external
impact, thus revealing linear relationship between logarithm of the coefficient
of thermodynamic activity and reaction shift from true thermodynamic
equilibrium in the state of open equilibrium. Discovered relationship prompts
us to use in many systems a combination of the linearity coefficient and
reaction shift from true equilibrium rather then activity coefficients. The
coefficient of linearity, and therefore coefficient of thermodynamic activity
can be found via thermodynamic computer simulation for systems with
known thermochemical properties. Numerical data obtained by various
simulation techniques proved premise of the developed in this work method
of chemical dynamics. Results of this work build up a good basis for new
approach and software for thermodynamic simulation of non-ideal chemical
and related systems.
INTRODUCTION: BACK TO CHEMICAL DYNAMICS.
“True”, or “internal” thermodynamic equilibrium is defined by current thermodynamic
paradigm only for isolated systems. That’s why applications of chemical thermodynamics
to real, mostly open, systems often lead to severe misinterpretation of their status, bringing
approximate rather than precise results. The subject of this work was to find a possibility
to expand the idea of equilibrium to open systems and a relationship between deviation of
a chemical system from “true” equilibrium and parameters of its non-ideality.
To explain differences between analytical and active concentrations of some components
in complex systems, conventional methods use excessive thermodynamic functions. It is
noteworthy that these functions and related coefficients of thermodynamic activity were
introduced rather for convenience [1], first being a fig leaf for the lack of our knowledge of
what’s going on in real systems.
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Among attempts to solve the problem one should mention works of D. Korzhinsky [2] and
then I. Karpov [3]. Considering interaction of open systems within the multisystem, the
method distinguishes between the common, or mobile components, and inert components,
which are specific for each subsystem and cannot be present in any other. The mobile
components, carrying intensive thermodynamic properties and thus contributing the
subsystem’s Gibb’s potential, are responsible for the subsystem’s interaction. Coefficients
of thermodynamic activity of the mobile components may vary while for the inert
components they do not have any physical sense [4].
In this work we used method De Donder, who had introduced the thermodynamic affinity,
interpreting it as a thermodynamic force, and reaction extent as a “chemical distance” [5].
We will use redefined reaction extent as dj=dnkj/kj (instead of original dj=dnkj/kj), or in
incrementsj =nkj/kj. Here nkj is the amount of moles, consumed or appeared in jreaction in its run between two arbitrary states, one of them usually is the initial state. The
kj value equals to a number of moles of k-component, consumed or appeared in an
isolated j-reaction on its way from initial state to true equilibrium and may be considered
a thermodynamic equivalent of chemical transformation. Redefined j is a dimensionless
chemical distance (“cd”) between initial and running states of j-reaction, 0=<j=<1, and
thermodynamic affinity A = - (G/)p,T turns into a classical force by definition,
customary in physics and related sciences (derivative of potential by coordinate).
Chemical reaction in isolated system is driven only by internal force, Aij,. True
thermodynamic equilibrium occurs at A´ij = 0, and at this point ´j = 1. Reactions in open
system are driven by both internal and external, Aej, forces [7] where the external force is
related to chemical or, in general, thermodynamic interaction of the open system with its
environment. If not far from “true” equilibrium, one can receive condition of the “open”
equilibrium from linear constitutional equations of non-equilibrium thermodynamics at
zero reaction rate with resultant affinity
A*ij + aie A*ej = 0,
(1)
where aie is the Onsager coefficient [6]. The accent mark and asterisk relate values to
isolated (“true”) or open equilibrium correspondingly.
In this work we will use only one assumption which in fact slightly extends the hypothesis
of linearity. We suppose that in the vicinity of true thermodynamic equilibrium the
reaction shift j =1 - j is directly proportional to the shifting force
j = ie Aej .
(2)
Recall that Ai =-(Gi /i ). Substituting (2) into (1), and retaining in writing only j for
j andj for j, we will have after a simple transformation
G*ij + bie *j *j = 0,
(3)
where bie = aie /ie. Corresponding constant equation is
G0ij + RTln*j (, *j) + bie * j *j = 0,
(4)
*j (, *j) is the activities product, mole fractions are expressed with reaction extent.
So, as soon as chemical system becomes open, the appropriate constant equations (and
system’s Gibbs’ potential) include a non-linear, non-classical term originated from the
interaction of the system with its environment.
Defining a new value - “non-thermodynamic”, or alternative temperature of open system
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Ta = bie / R,
(5)
where R is universal gas constant, we turn (4) to
G0 ij + RTt ln*j + RTa *j*j.
Recall well known classical expression G0 ij = - RTln Ki . Dividing (6) by (-RTt ),
presenting the activity product at open equilibrium as *j (kj, *j)= n0pj + pj
*j)/]pj/n0rj - rj, *j)/]rj and equilibrium constant as Ki = `kj,1) due to `j =1,
and defining reduced temperature as = Ta /Tt we transform equation (6) into
ln [`kj,1)/ kj , *j)] +  j*j*j = 0.
(7)
Being divided by *j , this equation expresses linearity between the thermodynamic force
and reaction shift in new terms
{ln [`kj,1)/ kj , *j)]}/*j = - j*j.
(8)
INVESTIGATION OF THE FORCE-SHIFT RELATIONSHIP.
First, consider the force expression from equation (8). Its numerator is a logarithm of a
combination of product molar parts for a given stoichiometric equation. The expression
under the logarithm sign is the molar parts product for ideal system divided by the same
product where kj replaced by a product (*jkj) due to the system’s shift from “true”
equilibrium. Table 1 represents functions `kj,1)/ kj , *j) for simple chemical
reactions with initial amounts of reactants A and B equal to one mole. Graphs of the
Table 1.
Thermodynamic forces for simple chemical reactions [left side of eq. (8)].
Reaction equation.
A + B = AB
A + 2B = AB2
2A + 2B = A2B2
{ln[`kj,1)/ kj , *j)]}/
[(2- )/(1-2-2)] / [(2-22)*(1-2+22)]
(1-
[(2-3 2-33 * [(1-24 * (1/
2
reaction shifts vs. thermodynamic forces are shown at Fig.1. One can see well expressed
linearity on the shift-force curves. Extent of the linearity along the shift axes depends on
the value.
Going down to real objects, consider a model system containing a double compound AR
and A relates to reactant A which belongs to the system (A,I) and is open to interaction
with R. Two competing processes take place in the system - decomposition of AR, or
control reaction (C): AR = A + R, and leading reaction (L): A + I = *L, the right side
in the last case represents a sum of products. Resulting reaction in the system is AR + I =
*L + R.
To obtain numbers for real substances, we used thermodynamic simulation (HSC
Chemistry for Windows) for a model set of substances. The Is were S, C, H2, and
MeORs were double oxides with symbol Me standing for Co, Ni, Fe, Sr, Ca, Pb and Mn.
As restricting parts R were used oxides of Si, Ti, Cr, and some others. Chosen double
compounds had relatively high negative standard change of Gibbs’ potential to provide
negligible dissociation in absence of I. In chosen systems the C-reactions were
(MeO)R=(MeO)+R, and L-reactions - (MeO)+I. Amount of the MeO moles consumed
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1
1
0,5
0,5
0,5
0
0
0
1
0 1 2 3 4 5
0 1 2 3 4 5
0 1 2 3 4 5
Fig. 1. Shift of some simple chemical reactions from true equilibrium (ordinate) vs.
shifting force (abscissa). Reactions, left to right, values of in brackets:
A+B=AB (, 0.9), A+2B=AB2 ( 2A+2B=A2B2
(.
in isolated (MeO+I) reaction between initial state and true equilibrium was taken as value
of kL. Reaction extents for open L-reactions with different Rs have been calculated as
quotients of consumed amounts of (MeO) (that is nkj) by kL. As numerator for the
thermodynamic force we used traditional GC (or even G0C which does not make a big
difference at moderate temperatures), and the force was equal to (-G0C /*L). Some of
the results for reactions (MeOR+S) are shown on Fig.2. In this group of reactions value of
(- G0C / *L) plays role of external thermodynamic force for (MeO+S) reaction.
 *
FeO*R
CoO*
0,5
CaO* R
Fext.
0
0
300
600
Fig.2. Shift *L vs. force (= - G0C / *L ), kJ/m·cd, 298.15K, direct thermodynamic
simulation. Points on the graphs correspond to various Rs.
The most important is the fact that in both cases the data, showing the reality of linear
relationship, have been received using exclusively current formalism of chemical
equilibrium where no such kind of relationship was ever assumed at all. It is quite
obvious that linear dependence took place in some cases up to essential values of
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deviation from equilibrium. Results shown on Fig.1 and Fig.2 prove the basics and some
conclusions of the method that we call a method of chemical dynamics for explicit usage
of thermodynamic forces.
FROM CHEMICAL DYNAMICS TO CHEMICAL THERMODYNAMICS:
THERMODYNAMIC ACTIVITY AND REACTION SHIFT AT OPEN EQUILIBRIUM.
In classical chemical thermodynamics excessive thermodynamic functions and
coefficients of thermodynamic activity are related by equation
Qj = - RTt ln kj,
(9)
and constant equation for non-ideal system with kj 1 is
G0j = - RTt ln *kj - RTtln*xkj.
For simplicity we omitted power values, equal to stoichiometric coefficients, and x kj are
molar fractions. The non-linear term of the equation (7) also belongs to a non-ideal
system, and comparison of (6) and (10) leads to following equality in open equilibrium
j *j ln * kj)/*j.
(11)

This result is quite understandable. For instance, in case of A R the chemical bond
between A and R reduces reaction activity of A; the same result will be obtained for
reaction (A + I) with reduced coefficient of thermodynamic activity of A.
In simple case of only one common component A the relationship between the L-shift
and activity coefficient of A is very simple
*L = (1/L)(-ln *)/* L].  (14)
where (-ln *)/*L] represents external thermodynamic force acting against L-reaction
and divided by RTt.. This expression for the force as well as the total equation (14) are
new. Equation (14) connects values from chemical dynamics with traditional values of
classical chemical thermodynamics. Yet again, at *L= 0 we have immediately *=1,
and vice versa, a correlation, providing an explicit and instant transition between open
and isolated systems. In case of multiple interactions one should expect additivity of the
shift increments, caused by interaction with different reaction subsystems, which
follows the additivity of appropriate logarithms of activity coefficients
Data on Fig. 3 were obtained using two different methods of thermodynamic simulation.
I-simulation relates to an isolated (AR+I) system with real R and A and AR=1 in all
cases. In O-simulation a combination of |A+Y2O3+I| represented the open system where
R was excluded and replaced by yttrium oxide, neutral to A and I) to keep the same total
amount of moles in the system as in I-simulation and avoid interaction between A and R.
Binding of A into double compounds with R, resulting in reduced reaction ability of A,
was simulated varying . I-simulation provided a relationship in corresponding rows of
the *L - * values, and O-simulation - with *L - G0A R correspondence. Standard
change of Gibbs’ potentialG0C, determining strength of the AR bond, was considered
an excessive thermodynamic function to the L-reaction.
We have calculated numeric values of L from the data used for plotting Fig. 3. For
complex oxides and simulation temperatures CoO*R/298, SrO*R/798, PbO*R/298,
corresponding values of L are 40.02, 6.54, and 3.93. Coefficients of determination are
higher than 0.97, and standard deviation falls within the range of (3 … 9)%.
111
P b O*
1
S rO*

0,5
Co O*
0
(-ln 
0
25

Fig. 3. *vs. (-ln *) (I-simulation, x) and vs. (G0A R/*) (O-simulation, o),
(MeOR+S). PbO and CoO at 298K, SrO at 798.
Strong relation between reaction shifts and activity coefficients means automatically
strong relation between shifts and excessive thermodynamic functions, and external
thermodynamic forces as well. Along with standard change of Gibbs’ potential we also
tried two others - QL which was calculated by equation (14) with *, used in the Osimulation, and another, G*L , found as a difference between G0L and equilibrium value
of RTtln*xiL. Referring to the same *L, all three should be equal or close in values,
and results were found to be very close.
One can immediately see strong linearity of the initial areas of the curves along with well
expressed break points, were the curves change their slope, turning from strong
dependence of the reaction shift upon external thermodynamic force to weak dependence.
This kind of “saturation” is particularly pronounced on the PbO* curve.
0.6
0.3
0.5
0.1
0.3
0
0
5
10
15
Fig. 4. *L vs. (-ln *). TDS. CoOR-C system, 298K, different reactant ratio.
Numbers at plots identify the initial value of the CoO/C ratio.
It is worthy to mention that open equilibrium may be defined using both external (like
G0C) and internal (the bound affinity, see [7]) values as well as, say, a neutral, or general
value like a function calculated by (14) at given activity coefficient. In principle all three
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may be used to calculate or evaluate L. That allows us to reword more explicitly the
problem, set in the beginning of this work and explain the alternative temperature more
clear. It is easy to see that equation (14) represents another form of the shift-force
linearity.
Because we ran thermodynamic simulation within a certain range of initial ratios between
components of the reaction mixtures it was interesting to see what a difference it made.
Fig. 4 shows no essential dependence of the L value on this ratio in the CoOR - C
system.
PRACTICAL APPLICATIONS.
We are unable to simulate and compute equilibrium composition of most complex
chemical systems if we don’t know appropriate coefficients of thermodynamic activity,
and their numeric values are very expensive. The method of chemical dynamics offers
easier and involving essentially less efforts way to run that kind of research. Indeed,
equilibrium of complex chemical system per this method may be interpreted as
equilibrium of subsystems’ shifting forces causing their deviation from “true” equilibrium
states. Having the kj values received from solution for the isolated state, we have equation
containing only *j as variable and j as a parameter of the theory, determining the
subsystem’s response to external thermodynamic impact
lnj =ln(kj*j) -j **j.

Current methods of simulation of complex equilibria use the constant equations (or
equivalent expressions if minimizing Gibbs’ potential of subsystems) in the same form as
if the subsystems are isolated. Their traditional joint solution is only to restrict consuming
the common participants and thus achieve material balance within the system thus playing
role of an accountant. Application of current methods to real systems leads to some errors
in simulation results originated due to misinterpretation of their status. Method developed
in this work treats states of subsystems as open equilibria within a complex equilibrium,
and leads to more correct numerical output. Joint solution to a system of equations (15)
allows us to simulate chemical equilibrium in an arbitrary complex system with higher
precision. Along with this opportunity, method of chemical dynamics provides us with a
simple way to calculate coefficients of thermodynamic activity following the above
described procedures.
A principal feature of application of the method consists in usage of reaction shift (as the
system’s response to external impact) multiplied by proportionality coefficient *j rather
than activity coefficient kj. Due to an easy way to obtain value of j by thermodynamic
simulation within minutes (not hours!), the method of chemical dynamics brings new
opportunities to analysis and simulation of complex chemical systems.
DISCUSSION.
The new basic equation received in this work links equilibrium and non-equilibrium
thermodynamics and may be rewritten more generally as
GG0j + RTt f t (*jRTa f a (*j.
The graph of the relationship between reaction shift and external force found in this work,
precisely repeats well known Hooke’s law [8] with its linearity at low elongations of a
stretched material and its yield point. In our case the yield point, where the curve sharply
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deviates from the straight line or, in some cases, just changes the slope, was very
distinctive on all plots. By analogy, the value of 1/Lmay be considered a coefficient, and
the yield point - a limit of thermodynamic proportionality of the chemical Hooke’s law.
We cannot tell to what extent the coefficient and the limit of thermodynamic
proportionality may be considered characteristics of a chemical elasticity thus providing
complete return of chemical reaction back to initial point when the chemical force returns
to zero value, i.e. with or without a sort of chemical hysteresis in force-shift coordinates.
This problem could be investigated in future research.
Addressing to a skeptical reader, we’d like to underline that all new results of this work
have been received and proven numerically within the current paradigm of chemical
thermodynamics. Our non-traditional term of the basic equation already existed in
chemical thermodynamics in form of excessive thermodynamic function. This work offers
alternative picture of its origin and its relation to an external impact on the chemical
system. From this point of view, we consider results of this work neither revolutionary nor
contradictory. We just tried to find out what has been lost or hidden when chemical
system, the major object of chemical thermodynamics, was idealized as an isolated entity.
One can find more detailed description of the method and its projection onto non-linear
thermodynamics in [9].
REFERENCES.
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Alberty, R. A.:
Physical Chemistry. New York, Wiley & Sons, 1983.
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Physico-Chimicheskie Osnovy Analiza Paragenezisa Mineralov (Physicochemical Basics of Analysis of Mineral Paragenesis). Moscow, AN USSR,
1956.
[3] Karpov, I. K.:
Physico-Chimicheskoje Modelirovanie na EVM v Geochimii (Physico-chemical
Computer Simulation in Geochemistry). Novosibirsk, Nauka, 1981.
[4] Zilbergleyt, B. J.: Russian J. Phys. Chem., 1985, 59(7), 1059.
[5] de Donder, T.; van Risselberge, T.:
Thermodynamic Theory of Affinity. Stanford, Oxford University Press, 1936.
[6] Gyarmati, I.:
Non-Equilibrium Thermodynamics. Berlin, Springer-Verlag, 1970.
[7] Zilbergleyt, B. J.: Russian J. Phys. Chem., 1985, 59(10), 1574.
[8] Parton, V. Z.; Perlin, P. I.:
Mathematical Methods in Theory of Elasticity. Moscow, MIR Publishers, 1984.
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http://arXiv.org/abs/physics/0004035, 2000, April 19.
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