Equilibrium-constant calculations involving gases quickly become more complicated than the ideas they were intended to illustrate, and we will defer most equilibrium calculations to the discussion of aqueous solutions in Chapter 16. However, it is useful to look briefly at the equilibrium constants for some of the reactions discussed previously in this chapter. Along the way we will encounter some of the fundamental ideas about manipulating equilibrium-constant expressions.

HCl Synthesis

The equilibrium-constant expression for the HCl reaction has the same form as that of HI synthesis:

(g) + (g) 2HCl (g)
G = -45.54 kcal per 2 moles of HCl




The experimental value of K, for this reaction is 2.5 X 10 . Notice that, with equal power of concentration terms in both numerator and denominator, for HCl is a unitless quantity.

for this reaction will have the same numerical value whether concentrations are measured in moles per liter, mole fractions, partial pressures, or any other convenient system ( = = = ). This is not always true.

If we begin with equal concentrations of and , then since they react in equal amounts, the concentrations of these two substances always will be equal, [] = [], and at equilibrium we can write:

or


This tells us that equilibrium will not be reached until the ratio of HCl to , (or ) has risen to 50 million billion to one! It is not surprising that we cannot detect any or in the products at equilibrium.

So large an equilibrium constant indicates that the reaction starting with equal concentrations of , , and HCl should be highly spontaneous, since the reaction has a long way to go before reaching equilibrium. The large negative standard free energy change, G= -45.54 kcal per two moles of HCl, indicates the same thing.