The Haber Process | ChemTalk (2024)

Core Concepts

In this article, you will learn about the Haber Process and its importance, as well as its kinetics, thermodynamics, and mechanisms.

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What is the Haber Process?

The Haber Process is an industrial method of reacting nitrogen and hydrogen gas to produce ammonia. Stoichiometrically, this involves one mole of diatomic nitrogen and three moles of diatomic hydrogen per two moles of ammonia.

Haber Process Equation: N₂ + 3H₂ → 2NH₃

Many chemists call the Haber Process a form of “nitrogen fixation,” meaning the conversion of nitrogen gas to a more chemically useful form. In nature, microorganisms called diazotrophs perform nitrogen fixation, which helps cycle nitrogen through ecological systems. Nitrogen, of course, serves as one of the fundamental elements of biological life, forming an important part of proteins and nucleic acids.

Nitrogen, in the form of ammonia, also has importance in research and industrial applications. However, despite diatomic nitrogen serving as the most abundant atmospheric gas, its lack of reactivity made it hard for chemists to harness.

In 1909, German chemist Fritz Haber made scientific history when he demonstrated synthesizing ammonia from atmospheric nitrogen and hydrogen. Later in 1911, the method was significantly improved by another German chemist Carl Bosch, which made the reaction more efficient for industry. As a result, when performed today, many refer to the method as the Haber-Bosch Process.

Due to the work of Haber and Bosch, ammonia can now be readily synthesized for a variety of important purposes, such as fertilizer, explosives, refrigeration technology, and pharmaceuticals.

But what exactly makes the Haber Process so efficient? Let’s take a closer look at the physical chemistry behind the reaction.

The Kinetics of the Haber Process

If nitrogen gas is so inert, how did Fritz Haber manage to get it to react? To get past the unreactivity of nitrogen, the Haber Process involves two important factors. The first factor is an effective catalyst.

Transition metals are almost always used to catalyze the Haber Process. Most often, industrial chemists use magnetite, an iron oxide (Fe₃O₄), with smaller amounts of potassium hydroxide (KOH). Other common catalysts include chromium and osmium oxides.

Mechanistically, the catalyst must bind to nitrogen gas. This allows hydrogen to add across the nitrogen bonds until the nitrogen fully saturates, yielding ammonia. Without this crucial binding of the catalyst, nitrogen remains inert and unreactive, producing virtually no ammonia.

The Thermodynamics of the Haber Process

With or without a catalyst, the reaction producing ammonia is spontaneous (∆G < 0). This is due to a fundamental aspect of physical chemistry: the kinetics and thermodynamics of a reaction remain often independent of each other. Though we may intuit that a highly-thermodynamically favored reaction occurs faster than one less favored, there exist exceptions. The Haber Process serves as one such exception; a reaction with favorable thermodynamics but incredibly slow kinetics.

Aside from the spontaneity, we can also observe the reaction has a negative change in entropy (∆S < 0). Specifically, we see four moles of reactant gas producing two moles of product gas.

N₂ + 3H₂ → 2NH₃

A reduction in the moles of gas indicates that our system has fewer possible microstates, indicating decreasing entropy. We can also justify the negative change by saying fewer gas molecules means more “disorder”.

For our reaction to be spontaneous despite the decreasing entropy, our reaction must have decreasing enthalpy. This is indeed the case: the Haber Process is exothermic (∆H < 0). According to the definition of Gibbs free energy, spontaneity must increase as temperature decreases. Put differently, decreasing temperature makes ∆S less significant, making ∆G more negative:

∆G = ∆H – T∆S

(In this equation, T indicates the temperature at which the reaction takes place)

This has important implications for the ideal temperature at which to perform the Haber Process, as we observe in the next section.

The Equilibrium Properties of the Haber Process

As we just covered, the Haber Process is exothermic and increases spontaneity at lower temperatures. With this information, chemists can increase the yield of the reaction by lowering temperatures. The explanation for this phenomenon involves the principles of chemical equilibrium.

Due to exothermicity, we could say that heat serves as a “product” of the reaction. According to L’Chatelier’s Principle, we can shift the equilibrium of a reaction mixture by placing stress on the system. The reaction then responds in the direction to counteract the stress. By cooling the mixture, we “remove” heat, and the system counteracts this stress by producing more products. Thus, we get more ammonia under lower temperatures.

However, if we decrease the temperature too much, the kinetics of our reaction decrease since the molecules don’t move as fast. To compromise, industrial chemists usually set the reaction temperature of the Haber Process at 450-500°C.

The Full Haber Process

Now that we know the physical chemistry, let’s take a look at the full process.

One thing you may notice is that dihydrogen gas is not directly inputted into the system. Instead, methane (1) and water (2) react to produce hydrogen within the system (3). Later, oxygen and nitrogen gas enter the system (4), which further generates hydrogen from methane (5). In both reactions, carbon monoxide forms as a by-product.

Further water vapor enters the system, oxidizing carbon monoxide to carbon dioxide (6). This carbon dioxide then leaves (8), producing a mixture of nitrogen and hydrogen gas.

This gas mixture becomes compressed to more than 200atm (7) and pre-heats to ~500°C (9). Then, the heated and pressurized mixture enters the reaction chamber, fitted with the transition metal catalyst (10). A certain amount of ammonia is produced, which is distilled by running a cold water jacket over the reaction mixture (11). This condenses the ammonia gas to a liquid, which drains out of the system (13).

Importantly, on the first run, only about 15% of the nitrogen and hydrogen will have reacted. The unreacted gas separates from the liquid ammonia (12) and enters the reaction chamber again. Eventually, about 98% of the original gas mixture reacts to form ammonia after enough runs.