Lab Manual: Homogeneous Chemical Equilibrium
Homogeneous Chemical Equilibrium
INTRODUCTION & THEORY
In beginning studies of chemistry, it is convenient to assume that in chemical reactions the reactant(s) are converted
entirely to product(s) provided that the moles undergoing reaction are in agreement with the stoichiometry of the balanced
When this is assumed or when the reaction is known to go very nearly to 100%, a one-headed arrow
separates reactants (A, B, and others) from products (G, H, and others) as in the balanced chemical equation:
In reality, many chemical reactions stop far short of 100% conversion of reactants to products.
It is even possible that very
little product is actually produced.
It may also be the case that it is impossible to produce any of the assumed products
since the reactions may be non-spontaneous. Both the spontaneity of a reaction (the direction in which it proceeds) and the
extent to which it proceeds are influenced by reaction conditions, such as temperature and pressure.
In addition, the
spontaneity of the reaction also depends upon the initial concentrations of reactants and products.
In this experiment, we
are interested in learning a very common mathematical method by which the extent to which a chemical reaction proceeds
toward product formation, or completion, can be expressed.
In all spontaneous chemical reactions the concentrations of the reactants decrease and the concentrations of the products
increase until a steady state condition is reached.
At this steady state or equilibrium condition, the concentrations of the
reactants and products no longer change with time.
As noted earlier, this condition of equilibrium may be established well short of 100% conversion of reactants to products.
The equilibrium state is denoted by using
instead of a one-headed arrow, in the balanced equation.
Since the concentrations of all species are constant at equilibrium, there exists a numerical ratio or constant, which
describes the relative concentrations of all species.
The general equation for the equilibrium constant, K
, in concentration
units, for the reaction,
is expressed as,
This is also known to as the Law of Mass Action.
It should be noted that the numerical magnitude of K
is a direct indication of the extent of completion of a chemical
reaction at equilibrium.
Large values of K
indicate high product and low reactant equilibrium concentrations and thus a
high degree of completion.
Small values of K
arise from very small equilibrium concentrations of product(s) and large
unreacted concentrations of reactant(s).
Many important reactions such as the one to be studied in this experiment have
values, somewhere between 0.001 and 100.
Reactions with moderate K
values are usually readily studied
because they will not result in very small concentrations of either the reactants or products.