Retro-Technics • author: a3310i • last modified: 2021.12.22 •

Nitrogen (Latin name nitrogenium), abbreviated N, is an element with an atomic weight of 14.008. Under normal conditions, it is a colourless and odourless gas. The atomic number is 7. Nitrogen therefore has five electrons in its last orbit. It is still three electrons short of the most stable arrangement, the octet. Consequently, nitrogen gives up electrons - and its maximum valence is 5. This element was probably discovered in 1777 by a Swedish scientist named Scheele.
How much nitrogen is there in nature? Nitrogen is a component of proteins found in living organisms. However, there is very little of it in this form. Slightly more nitrogen is found in the earth's crust. It is mainly found in the form of sodium and potassium nitrates. These compounds are known as saltpetre, and there are rich deposits in central Asia and South America. The most abundant source of nitrogen, however, is the earth's atmosphere. On average, 100 ml of atmospheric air contains 78 ml of nitrogen, or by weight, there is 75.4 g of nitrogen in 100 g of air. However, only a few organisms are able to take advantage of the abundance of atmospheric nitrogen.

[001]

Today, a developed nitrogen fertiliser industry is the basis for high agricultural productivity. In the mid-19^th century, when nitrogen fertilisers were not yet used, European oat yields did not exceed 0.56 tonnes per hectare. In the 20^th century, however, artificial fertilisers, mainly nitrogenous, have resulted in yields of up to 3 tonnes per hectare.
One of the nitrogen compounds that was probably the earliest known to man was nitrate. Already in the past, chemists were able to produce nitric acid by the action of sulphuric acid on sodium nitrate. This acid was highly valued, because its mixture with hydrochloric acid allowed the dissolution of gold. With time, when the idea of transmutation and the philosopher's stone fell into disuse, nitric acid lost its alchemical significance and began to find practical applications.
The second nitrogenous compound, also known to alchemists, was the so-called spirit of deer horn. It is difficult for a person not familiar with the old nomenclature to guess that under this allegorical name is hidden a simple and popular compound - ammonia. The old name probably has something to do with the dry distillation of animal waste such as horns, hooves and leather. During this process, proteins are broken down and nitrogen is released in the form of ammonia.
The third nitrogenous compound long known to man was potassium nitrate, KNO₃.

The demand for nitric acid rose sharply only in the middle of the 19^th century. Here, in 1845, the Swiss chemist Schonbein, examining cotton, found that by acting on it with sulphuric and nitric acid it could be transformed into an extraordinarily strong and violently explosive material. This product was called gun cotton. However, being aware of the destructive effect it could have, he kept his discovery a secret. Despite the best intentions of the chemist, the news about guncotton leaked out from his workshop and soon in Germany factories of this material were established, working, of course, for the army. A massive explosion at one such factory in 1848 destroyed almost the entire town. Hundreds of casualties and millions in material losses, however, did not deter more and more factory owners from producing gun cotton.
After gun cotton came the invention of an even more powerful explosive, nitroglycerine. Both discoveries immediately became popular - especially when gun cotton was used to manufacture smokeless gunpowder and nitroglycerine was used to create dynamite. There was no shortage of wars, and firearms technology advanced rapidly with a panache worthy of a better cause. No wonder that in this situation sulphuric and nitric acid were promoted to the dignity of strategic intermediates. Whereas sulphuric acid could be produced from raw materials found in almost every country, the situation with nitric acid was much worse. Europe did not have any deposits of sodium nitrate, so this compound began to be imported from distant Chile. It is known that little has changed over the centuries, because until the first years of the 20^th century, nitric acid was still produced using a method developed by alchemists. Apart from the primitiveness of this method and the huge consumption of sulphuric acid, there was complete dependence on importing the raw material, Chilean nitrate.

[002] Gun cotton.

In view of these inconveniences, the matter centred on the question of how to bind atmospheric nitrogen. It was originally assumed that this would be done by synthesising nitric oxide. However, this was found to be almost impossible. The reaction:

is an extremely endothermic reaction and requires 21.6 kcal of energy per mole of product. After numerous research works, a solution to the problem of nitric oxide synthesis has finally been found, at least theoretically - the synthesis should be carried out at a temperature of at least 3000 °C.
Among the constructors of arc furnaces for the synthesis of nitric oxide, the name of the Polish physicist and chemist Professor Ignacy Mościcki came to the fore. In order to increase the surface area of the arc, he placed it in a strong magnetic field. In turn, by selecting the appropriate aerodynamic conditions, such as the shape of the vents and the speed of the gas flow, he achieved that the arc ignited between the electrodes spun in a kind of fiery disc with a large diameter. A very strong stream of air flowed through the disc. After passing through the hot zone, where the synthesis took place, the gases were immediately directed to a strongly cooled element, where they were cooled down. Further oxidation of the resulting nitric oxide to nitrogen dioxide:

and then dissolving it in water was no longer a problem.
The most serious disadvantage of the arc method was the consumption of large amounts of electricity. On average, more than 60 kWh of electricity had to be used to bind 1 kg of nitrogen from the air. No wonder that only countries with cheap electricity could afford to set up nitric acid factories of this type. Despite the undisputed victory of science, the arc method was not economical. Many chemists and designers looked for a cheaper method of fixing atmospheric nitrogen. Two German chemists, Haber and Bosch, had excellent success in this field. Haber, knowing the results of his work on the direct oxidation of nitrogen, decided to obtain nitric acid not from the air but by a somewhat circuitous route, via ammonia. Why ammonia? Firstly, the hydrogen-nitrogen bonding reaction is slightly exothermic, in contrast to the oxide synthesis, where a lot of energy had to be supplied. The second reason was raw material considerations. Namely, hydrogen is a waste gas in many chemical plants. The last consideration is that nitric acid can easily be obtained from ammonia. In developing the synthesis of ammonia, however, chemists encountered even greater difficulties than in the synthesis of nitric oxide in arc furnaces. The reaction for the synthesis of ammonia follows the equation:

The two-directional arrow separating the raw and the substrate show that the reaction can proceed either to the left or to the right. Which direction the reaction moves in depends on external factors. In order to better understand the role of the determining factors, it is necessary to know Le Chatelier's rule. This rule states that if a system in equilibrium is subjected to an external action, then such a transformation takes place in it that partially abolishes this action.

For instance, if the solubility of a body in water increases with temperature, then the solution cools down as it dissolves. In 100g of H₂O at 0 °C 13 g of potassium nitrate dissolves, at 100 °C over 245 g. These data reveal that the solubility of potassium nitrate increases markedly with temperature. However, if a few teaspoons of potassium nitrate are poured into a glass of water, the solution cools down as it dissolves. This will be precisely the effect resulting from Le Chatelier's rule. It also works the other way round - if the solubility of a compound decreases as the temperature increases, then as such a body dissolves, the solution heats up. Thus, compounds that require a large amount of energy to form at high temperatures are stable. Conversely, compounds formed as a result of exothermic reactions can be stable only at low temperatures. This was the case with ammonia.
In the equation N2 + 3H₂ ↔2NH₃, you can see that there are four gas volume units on the left-hand side of the equation and only two on the right-hand side. According to Le Chatelier's principle, the resulting ammonia will decompose to counteract the decrease in volume. The problem of volume was solved by carrying out the synthesis under considerable pressure. If pressure is applied to the reacting gases, the system will try to reduce this pressure. This will happen by nitrogen combining with hydrogen, because then the volume decreases twice. In other words, the higher the pressure, the better the efficiency of ammonia synthase. All this has not yet solved the problem of temperature. At what temperature should the synthesis be carried out? A high temperature risks decomposition, while a low temperature means a slow reaction rate.

[003] Haber and Bosch method. Technology for the high-pressure synthesis of ammonia. 1. oil filter 2. reactor 3. compressor 4. circulating pump 5. boiler (optional) 6. water cooler 7. heat exchanger 8. filter 9. ammonia cooler 10. separator 11. exhaust gas stream 12. ammonia

After analysing all the parameters of the process, the German chemist decided that the role of the catalyst was essential. The catalyst would accelerate the reaction at a low temperature. After thousands of trials and hard work with varying success, Haber selected a set of catalysts that provided good performance while keeping the process temperature low. The development of this technology was a great achievement for science and a triumph for catalysis. The work was also pioneering, as catalyst science was just beginning to develop in the early 20^th century.