These values for , raise some paradoxical questions. At 2300K inside the combustion chamber, appreciable amounts of NO will be formed from and from the intake of air.

Example. At equilibrium a 1-liter steel tank at 2300K contains 0.04 mole of , 0.01 mole of , and an unknown amount of NO. How many moles of NO are present, and what is the mole fraction of NO?

Solution. The equilibrium-constant expression for the reaction is:




[NO] = 0.00081 mole liter
total moles = 0.04 + 0.01 + 0.0008 = 0.0508

= 0.0008 / 0.0508 = 0.016

The hot engine gases thus are 1.6% (mole percent) NO, which is an appreciable amount. However, if we repeat the calculations of the preceding example, but at room temperature, we arrive at the paradoxical conclusion that NO should be no problem in air pollution; equilibrium at 298K fies far on the side of and :

= 2.20 x 10

[NO] = 5.2 x 10 mole liter within the steel tank.

This quantity of NO would be totally undetectable and unimportant.

Where is the flaw in the analysis? The oxide NO obviously is a serious problem in photochemical smog. It is produced at high temperatures in internal combustion engines, but why does it not break down spontaneously into and as the gases rush out the tailpipe and are cooled down?

The problem is one of rates of reaction, and not of equilibrium. True, NO at equilibrium at 298K should break down into and , but the gases in our atmosphere are far from equilibrium. The breakdown reaction is a very slow one, and NO remains intact long enough to be oxidized to and enter the smog-producing pathway. Rates of reaction are important, and "How fast?" often is a more important question than "How far?".

"How fast?" is the question to be considered in the following chapter.