# 31.10: The Haber-Bosch Reaction Can Be Surface Catalyzed


## The Haber-Bosch Process for Synthesis of Ammonia

An example of an industrial catalytic process is the Haber-Bosch process. Karl Bosch (1874–1940) was a German chemical engineer who was responsible for designing the process that took advantage of Fritz Haber’s discoveries regarding the N2 + H2/NH3 equilibrium to make ammonia synthesis via this route cost-effective. He received the Nobel Prize in Chemistry in 1931 for his work. The industrial process is called either the Haber process or the Haber-Bosch process. used to synthesize ammonia via the following reaction:

$N_{2(g)}+3H_{2(g)} \rightleftharpoons 2NH_{3(g)} \nonumber$

with

$ΔH_{rxn}=−91.8\; kJ/mol \nonumber$

Because the reaction converts 4 mol of gaseous reactants to only 2 mol of gaseous product, Le Chatelier’s principle predicts that the formation of NH3 will be favored when the pressure is increased. The reaction is exothermic, however (ΔHrxn = −91.8 kJ/mol), so the equilibrium constant decreases with increasing temperature, which causes an equilibrium mixture to contain only relatively small amounts of ammonia at high temperatures ($$\PageIndex{1}$$). Taken together, these considerations suggest that the maximum yield of NH3 will be obtained if the reaction is carried out at as low a temperature and as high a pressure as possible. Unfortunately, at temperatures less than approximately 300°C, where the equilibrium yield of ammonia would be relatively high, the reaction is too slow to be of any commercial use. The industrial process, therefore, uses a mixed oxide (Fe2O3/K2O) catalyst that enables the reaction to proceed at a significant rate at temperatures of 400°C–530°C, where the formation of ammonia is less unfavorable than at higher temperatures.

Because of the low value of the equilibrium constant at high temperatures (e.g., K = 0.039 at 800 K), there is no way to produce an equilibrium mixture that contains large proportions of ammonia at high temperatures. We can, however, control the temperature and the pressure while using a catalyst to convert a fraction of the N2 and H2 in the reaction mixture to NH3, as is done in the Haber-Bosch process. This process also makes use of the fact that the product—ammonia—is less volatile than the reactants. Because NH3 is a liquid at room temperature at pressures greater than 10 atm, cooling the reaction mixture causes NH3 to condense from the vapor as liquid ammonia, which is easily separated from unreacted N2 and H2. The unreacted gases are recycled until the complete conversion of hydrogen and nitrogen to ammonia is eventually achieved. Figure $$\PageIndex{2}$$ is a simplified layout of a Haber-Bosch process plant.

We can write the seven reactions that are involved in the process where ad indicates that the molecule or atom is adsorbed on the catalyst

$H_{2}(g) + 2 S(s) \rightleftharpoons 2H \mathit{(ad)} \label{1}$

$N_{2}(g) \rightleftharpoons N_2 \mathit{(ad)} \label{2}$

$N_2 \mathit{(ad)} + 2 S(s) \rightleftharpoons 2N \mathit{(ad)} \label{3}$

$H \mathit{(ad)} + N \mathit{(ad)} \rightleftharpoons NH \mathit{(ad)} \label{4}$

$NH \mathit{(ad)} + H \mathit{(ad)} \rightleftharpoons NH_2 \mathit{(ad)} \label{5}$

$NH_2 \mathit{(ad)} + H \mathit{(ad)} \rightleftharpoons NH_3 \mathit{(ad)} \label{6}$

$NH_3 \mathit{(ad)} \rightleftharpoons NH_{3}(g) \label{7}$

Reaction $$\ref{3}$$ is much the slowest, and therefore the rate determining step. Figure $$\PageIndex{3}$$ summarizes the reaction scheme.

Gerhard Ertl worked out the energetics of the reaction shown in Figure $$\PageIndex{4}$$ below shows the amount of energy per mole needed on the catalyst and that which would be needed in the gas phase. Ertl's Nobel Prize speech about his work on the catalytic reactions forming ammonia and other catalytic reactions can be viewed on line.

Further studies of the Fe2O3/K2O catalyst have shown that the rate of the reaction depends on the particular surface on which the reaction is occurring. Figure $$\PageIndex{5}$$ shows the reaction rates for the synthesis of ammonia on five different surfaces of the iron.

### Contributors

• Anonymous

Modified by Joshua Halpern (Howard University), Scott Sinex, and Scott Johnson (PGCC)

Figure $$\PageIndex{3}$$ is from the ESA, now available on the internet wayback machine

Figure $$\PageIndex{4}$$ is from the Wikipedia Commons

31.10: The Haber-Bosch Reaction Can Be Surface Catalyzed is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts.