# Question #743d0

Oct 23, 2015

True. Haber process represents the application of LeChâtlier's Principle. It also represents the use of catalyst to reduce the activation energy.

#### Explanation:

${N}_{2} \left(g\right) + 3 {H}_{2} \left(g\right) r i g h t \le f t h a r p \infty n s 2 N {H}_{3} \left(g\right) \ldots e q n \left(1\right)$
$\Delta {H}_{{300}^{o}} = - 92.0 \frac{k J}{m o l}$

The rate of reaction and yield for eqn (1) are typically small to produce economical quantity of $N {H}_{3}$ and thus, it is desirable to increase the rate of reaction and yield. While increasing temperature can achieve higher rate of reaction, it is not desirable in this case as the reaction is an exothermic reaction (i.e. the bond formation of $N {H}_{3}$ release more energy than the energy required for bond breaking of ${N}_{2}$ and ${H}_{2}$). This means that increase of temperature will provide heat and thereby suppress bond forming activity (hence lower yield).

One solution to increase rate of reaction is through the reduction of activation energy (i.e. reduce energy required during bond breaking) using catalyst. Another solution is through the application of LeChâtlier's Principle. Since each gaseous molecule contribute to pressure (see Ideal Gas Law), the application of pressure will force the reaction to proceed in the direction with the lower number of gaseous molecule. In this particular case, 4 gaseous reactant produces 2 gaseous molecule of $N {H}_{3}$ . Hence, the equilibrium favors the production of $N {H}_{3}$ thereby increasing yield.

References:-
[1] The Synthetic Nitrogen Industry in World War I: Its Emergence and Expansion By Anthony S. Travis (pg 50)
[2] Prof. Shakhashir, Chemical of the Week,
http://scifun.chem.wisc.edu/chemweek/pdf/ammonia.pdf