Ammonia is manufactured on the industrial scale by the Haber process using the direct reaction of nitrogen with hydrogen at high pressure (10² – 10³ atm) and high temperature (400 – 550 °C) over a catalyst (e.g., α-iron), Equation 1.

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On the smaller scale ammonia is prepared by the reaction of an ammonium salt with a base, Equation 2, or hydrolysis of a nitride, Equation 3. The latter is a convenient route to ND₃ by the use of D₂O.

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The nitrogen in ammonia adopts sp³ hybridization, and ammonia has an umbrella structure (Figure 1) due to the stereochemically active lone pair.

Figure 1: The structure of ammonia.

The barrier to inversion of the umbrella is very low (E_a = 24 kJ/mol) and the inversion occurs 100’s of times a second, Equation 4. As a consequence it is not possible to isolate chiral amines in the same manner that is possible for phosphines.

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In a similar manner to water, Equation 5, ammonia is a self-ionizing, Equation 5; however, the equilibrium constant (K = 10^-33) is much lower than water (K = 10^-14). The lower dielectric constant of ammonia (16.5 @ 20 °C) as compared to water (80.4 @ 20 °C) means that ammonia is not as good as water as a solvent for ionic compounds, but is better for covalent organic compounds.

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The similarity of ammonia and water means that the two compounds are miscible. In fact, ammonia forms a series of solid hydrates, analogous to ice in which hydrogen bonding defines the structures (Figure 2). Several hydrates of ammonia are known, including: NH₃.2H₂O (ammonia dihydrate, ADH), NH₃.H₂O (ammonia monohydrate, AMH), and 2NH₃.2H₂O (ammonia hemihydrate, AHH).

Figure 2: The crystal structure of ammonia monohydrate (ANH-II) with hydrogen atoms omitted. Adapted from A. D. Fortes, E. Suard, M. -H. Lemée-Cailleau, C. J. Pickard, and R. J. Needs, J. Am. Chem. Soc., 2009, 131, 13508. Copyright: American Chemical Society (2009).

It should be noted that these hydrates do not contain discrete NH₄⁺ or OH^- ions, indicating that ammonium hydroxide does not exist as a discrete species despite the common useage of the name. In aqueous solution, ammonia is a weak base (pK_b = 4.75), Equation 7.

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Ammonia is a Lewis base and readily forms Lewis acid-base complexes with both transition metals, Equation 8, and main group metals (Figure 3).

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Figure 3: The molecular structure of AlMe₂(BHT)NH₃. Hydrogen atoms, except those attached to nitrogen, are omitted for clarity. Adapted from M. D. Healy, J. T. Leman, and A. R. Barron, J. Am. Chem. Soc., 1991, 113, 2776. Copyright: American Chemical Society.

The formation of stable ammonia complexes is the basis of a simple but effective method of detection: Nessler’s reagent, Equation 9. Using a 0.09 mol/L solution of potassium tetraiodomercurate(II), K₂[HgI₄], in 2.5 mol/L potassium hydroxide. A yellow coloration indicates the presence of ammonia: at higher concentrations, a brown precipitate may form. The sensitivity as a spot test is about 0.3 μg NH₃ in 2 μL.

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Ammonia forms a blue solution with Group 1 metals. As an example, the dissolution of sodium in liquid ammonia results in the formation of solvated Na⁺ cations and electrons, Equation 10 where solv = NH₃. The solvated electrons are stable in liquid ammonia and form a complex: [e^-(NH₃)₆].

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It is this solvated electron that gives the strong reducing properties of the solution as well as the characteristic signal in the ESR spectrum associated with a single unpaired electron. The blue color of the solution is often ascribed to these solvated electrons; however, their absorption is in the far infra-red region of the spectrum. A second species, Na^-(solv), is actually responsible for the blue color of the solution.

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The reaction of ammonia with oxygen is highly favored, Equation 12, and the flammability limit of ammonia is 16 – 25 vol%. If the reaction is carried out in the presence of a catalyst (Pt or Pd) the reaction can be limited to the formation of nitric oxide (NO), Equation 13.

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The ammonium cation (NH₄⁺) behaves in a similar manner to the Group 1 metal ions. The solubility and structure of ammonium salts particularly resembles those of potassium and rubidium because of their relative size (Table 1). One difference is that ammonium salts often decompose upon heating, Equation 14.

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The decomposition of ammonium salts of oxidizing acids can often be violent to highly explosive, and they should be treated with care. For example, while ammonium dichromate, (NH₄)₂Cr₂O₇, decomposes to give a volcano (Figure 4), ammonium permanganate, NH₄[MnO₄], is friction sensitive and explodes at 60 °C. Ammonium nitrate, NH₄[NO₃], can cause fire if contacted with a combustible material and is a common ingredient in explosives since it acts as the oxygen source due to its positive oxygen balance, i.e., the compound liberates oxygen surplus to its own needs upon decomposition, Equation 15.

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Figure 4: A laboratory demonstration of an ammonium dichromate volcano.