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Далеко занимљивији и важнији су утицаји промене фактора под којима се реакција одвија ([[температура]], [[притисак]], концентрације...) на равнотежу реакције. Претпоставке у вези са овим даје [[Ле Шатељеов принцип]]. Он има велике практичне импликације.
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==Chemical equilibrium in solution/gas phase reactions==
For illustration, consider the generic reversible reaction in solution
(or in the gas phase)
:<math>
mA + nB \leftrightarrow pC + qD
</math>
By the law of [[mass action]], the forward rate should equal
<math>k_{AB} [A]^{m} [B]^{n}</math>, whereas the backward rate
should equal <math>k_{CD} [C]^{p} [D]^{q}</math>, where
<math>k_{AB}</math> and <math>k_{CD}</math> are the forward
and backward [[reaction rate constant]]s, respectively,
and <math>[A], [B]</math>, etc. represent the concentrations
(or, more correctly, the
[[activity (chemistry)|chemical activities]]) of the reactants
and products. Setting the forward
and backward rate constants equal and cross-dividing, we arrive
at the [[equilibrium constant]] <math>K_{eq}</math>
:<math>
K_{eq} \equiv \frac{k_{AB}}{k_{CD}} =
\frac{\left[C\right]^{p} \left[D\right]^{q}} {\left[A\right]^{m} \left[B\right]^{n}}
</math>
Of course, anyone could prepare a solution in which the ratio of
concentrations on the right-hand side of the equation (called the
[[reaction quotient]]) did not equal <math>K_{eq}</math>. In such
a solution, the concentrations would not be at equilibrium; they would
start changing until the ratio of concentrations ''did'' equal
<math>K_{eq}</math>. Thus, the concentrations in this system are at
equilibrium (i.e., don't change with time) only if the reaction
quotient equals <math>K_{eq}</math>, and vice versa.
Suddenly adding more reactant (say, [A]) to a system in equilibrium
drives the equilibrium to the right (i.e., towards higher [C] and [D] concentrations and lower [B]). The sudden addition of [A] increases the
instantaneous ''forward'' rate without changing the ''backward'' rate.
Thus, the addition of [A] will cause C and D to be made faster and B to
be lost faster than the reverse reactions. Eventually, the system will
reach a new equilibrium where the ratio of concentrations exactly
equals <math>K_{eq}</math>.
The equilibrium position of a [[reaction]] is said to lie ''far to the right'' if, at equilibrium, nearly all the reactants are used up and ''far to the left'' if hardly any product is formed from the reactants. Changing the conditions of a reaction can shift the equilibrium to the right or left.
The kinetics of a reaction can be changed without altering its equilibrium concentrations. Specifically, the forward and backward rate constants can be both multiplied by the same factor without affecting their ratio (the equilibrium constant). This situation occurs quite commonly when a [[catalyst]] (such as an [[enzyme]]) is added to a reaction. Thus, the same equilibrium constant can be found in very fast and very slow reactions, and a fast forward reaction (by itself) does not imply that the reaction equilibrium lies far to the right.
In solids or other situations, the forward and backward rates may be
described by different equations, but one can usually define an
equivalent equilibrium constant by equating the forward and
backward rates and factoring out the constants (such as <math>k_{AB}</math>
and <math>k_{CD}</math>) from the variables (such as [A] and [B]).
==Chemical equilibrium and thermodynamics==
Although chemical equilibrium is defined kinetically (forward and
backward rates are equal), its properties can be studied thermodynamically,
i.e., from the free energies of the reactants and products. The main
equation is
:<math>
\Delta G^\circ = -RT \ln K_{eq}
</math>
where ΔG° is the standard free energy difference
between the products and reactants (e.g., in kcal/mol), <math>T</math> is
the absolute temperature in [[kelvin]]s and <math>R</math> is the universal
gas constant. This equation can be written equivalently as
:<math>
K_{eq} = e^{-\frac{\Delta G^{\circ}}{RT}}
</math>
Thus, the equilibrium constant depends on the temperature <math>T</math>
and also on variables that affect ΔG°, such as
temperature, pH, other co-solvents, etc.
==Examples of chemical equilibrium==
A common example given is the [[Haber process|Haber-Bosch process]], in which [[hydrogen]] and [[nitrogen]] combine to form [[ammonia]]. Equilibrium is reached when the rate of production of ammonia equals its rate of decomposition. [[Le Chatelier's principle]] describes qualitative predictions that can be made about a chemical equilibrium.
Classical equilibria are that between the colorless [[nitrogen dioxide]] and the brown [[dinitrogen tetroxide]] and the [[Schlenk equilibrium]].
In practice, most sets of reversible reactions have a stable
equilibrium. In rare cases, the concentrations may not settle to
fixed equilibrium values, but rather oscillate indefinitely.
==References==
<references />
2. P W Atkins Physical Chemsitry, Oxford University Press.
==Види још==
*[[Mass action]]
*[[Equilibrium constant]]
*[[Acidity constant]]
*[[Solubility constant]]
*[[Mass balance]]
*[[Free energy]]
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