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as their Reverse Reaction s. The rates of the forward and reverse reactions are generally not zero but, being equal, there are no net
changes in any of the reactant or product concentrations. This process is known as Dynamic equilibrium Atkins, Peter; Jones, Loretta. ''Princípios de química : Questionando a vida moderna e o meio ambiente.'' Tradução por Ignez Caracelli et alii. Porto Alegre : Bookman, 2001. (Translated from Atkins, Peter; Jones, Loretta. ''Chemistry: the quest for insight''). Vaibhav Patel 2005, Christian Hart 2006, Glen Paxman 2006. Leventhorpe publications.


CHEMICAL EQUILIBRIUM IN SOLUTION/GAS PHASE REACTIONS


For illustration, consider the generic reversible reaction in solution
(or in the gas phase)

:
mA + nB \leftrightarrow pC + qD


By the law of Mass Action , the forward rate should equal
k_{AB} [B ^{n}, whereas the backward rate
should equal k_{CD} [D ^{q}, where
k_{AB} and k_{CD} are the forward
and backward Reaction Rate Constant s, respectively,
and [B , etc. represent the concentrations
(or, more correctly, the
Chemical Activities ) of the reactants
and products. Setting the forward
and backward rate constants equal and cross-dividing, we arrive
at the Equilibrium Constant K_{eq}

:
K_{eq} \equiv rac{k_{AB}}{k_{CD}} =
rac{\left[C ight]^{p} \left[D ight]^{q}} {\left[A ight]^{m} \left[B ight]^{n}}


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 K_{eq}. In such
a solution, the concentrations would not be at equilibrium; they would
start changing until the ratio of concentrations ''did'' equal
K_{eq}. Thus, the concentrations in this system are at
equilibrium (i.e., don't change with time) only if the reaction
quotient equals K_{eq}, and vice versa.

Suddenly adding more reactant (say, {Link without Title} ) to a system in equilibrium
drives the equilibrium to the right (i.e., towards higher 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 {Link without Title} 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 K_{eq}.

If the forward and backward rate constants are both multiplied by
the same factor, the forward and backward rates both change but
the ratio (the equilibrium constant) doesn't change. This situation
occurs quite commonly when a Catalyst (such as an Enzyme ) is
added to a reaction. It's very important to understand that the
kinetics of a reaction can be changed without affecting its equilibrium.

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 k_{AB}
and k_{CD}) from the variables (such as and [B ).

The equilibrium position of a Reaction is said to lie ''far to the right'' if 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 result in a shift to the right or to the left of the equilibrium position.


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

:
\Delta G^\circ = -RT \ln K_{eq}


where \Delta G^\circ is the standard free energy difference
between the products and reactants (e.g., in kcal/mol), T is
the absolute temperature in Kelvin and R is the universal
gas constant. This equation can be written equivalently as

:
K_{eq} = e^{- rac{\Delta G^{\circ}}{RT}}


Thus, the equilibrium constant depends on the temperature T
and also on variables that affect \Delta G^{\circ}, such as
temperature, pH, other co-solvents, etc.


EXAMPLES OF CHEMICAL EQUILIBRIUM


A common example given is the 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.


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