Exp 2 - Lab Manual: Homogeneous Chemical Equilibrium 1...

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Lab Manual: Homogeneous Chemical Equilibrium 1 Homogeneous Chemical Equilibrium Buret - http://www.dartmouth.edu/~chemlab/techniques/titration.html Pipet - http://www.dartmouth.edu/~chemlab/techniques/pipet.html 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 chemical equation. 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: ... ... aA bB gG hH   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 c , in concentration units, for the reaction, ... ... aA bB gG hH   is expressed as, [] [ ] . . . . . . gh c ab GH K AB This is also known to as the Law of Mass Action. It should be noted that the numerical magnitude of K c is a direct indication of the extent of completion of a chemical reaction at equilibrium. Large values of K c indicate high product and low reactant equilibrium concentrations and thus a high degree of completion. Small values of K c 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 moderate K c values, somewhere between 0.001 and 100. Reactions with moderate K c values are usually readily studied because they will not result in very small concentrations of either the reactants or products.
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This note was uploaded on 06/02/2011 for the course CHEM 100A taught by Professor Dai during the Winter '06 term at UCSD.

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Exp 2 - Lab Manual: Homogeneous Chemical Equilibrium 1...

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