Nitrate

Nitrate is a polyatomic ion with the chemical formula NO3−. Salts containing this ion are called nitrates. Nitrates are common components of fertilizers and explosives. Almost all inorganic nitrates are soluble in water. An example of an insoluble (inorganic) nitrate is bismuth oxynitrate. In nature, nitrates are produced by a number of species of nitrifying bacteria in the natural environment using ammonia or urea as a source of nitrogen and source of free energy. Nitrate compounds for gunpowder were historically produced, in the absence of mineral nitrate sources, by means of various fermentation processes using urine and dung. Modern nitrate production is mostly focused on creation for fertilizer and chemical manufacturing for various applications, such as medicine synthesis, ceramics and preservation of meat. Annually, about 195 million metric tons of synthetic nitrogen fertilizers are used worldwide, with nitrates constituting a significant portion of this amount. Because nitrates are soluble and easily can be swept away from the soil because of precipitation, excessive agricultural use has been associated with nutrient runoff, water pollution, and the proliferation of aquatic dead zones. Direct exposure of nitrates for humans can have direct health consequences: the excess consumption of nitrates in cured meats is associated with intestinal cancers.

Chemical structure

The nitrate anion is the conjugate base of nitric acid, consisting of one central nitrogen atom surrounded by three identically bonded oxygen atoms in a trigonal planar arrangement. The nitrate ion carries a formal charge of −1. This charge results from a combination formal charge in which each of the three oxygens carries a −2⁄3 charge, whereas the nitrogen carries a +1 charge, all these adding up to formal charge of the polyatomic nitrate ion. This arrangement is commonly used as an example of resonance. Like the isoelectronic carbonate ion, the nitrate ion can be represented by three resonance structures:

Chemical and biochemical properties In the NO3− anion, the oxidation state of the central nitrogen atom is V (+5). This corresponds to the highest possible oxidation number of nitrogen. Nitrate is a potentially powerful oxidizer as evidenced by its explosive behaviour at high temperature when it is detonated in ammonium nitrate (NH4NO3), or black powder, ignited by the shock wave of a primary explosive. In contrast to red fuming nitric acid (HNO3/N2O4), or concentrated nitric acid (HNO3), nitrate in aqueous solution at neutral or high pH is only a weak oxidizing agent in redox reactions in which the reductant does not produce hydrogen ions (such as mercury going to calomel). However, it is still a strong oxidizer when the reductant does produce hydrogen ions, such as in the oxidation of hydrogen itself. Nitrate is stable in the absence of microorganisms, or reductants such as organic matter. In fact, nitrogen gas is thermodynamically stable in the presence of 1 atm of oxygen only in very acidic conditions, and otherwise would combine with it to form nitrate. This is shown by subtracting the two oxidation reactions:

N2 + 6 H2O → 2 NO3− + 12 H+ + 10 e− E 0 = 1.246 − 0.0709 pH + 0.0591 10 log ⁡ ( N O 3 − ) 2 P N 2 {\displaystyle \qquad E_{0}=1.246-0.0709{\text{ pH }}+{\frac {0.0591}{10}}\log {\frac {(\mathrm {NO_{3}^{-}} )^{2}}{P_{\mathrm {N_{2}} }}}}

2 H2O → O2 + 4 H+ + 4 e− E 0 = 1.228 − 0.0591 pH + 0.0591 4 log ⁡ P O 2 {\displaystyle \qquad \qquad \qquad E_{0}=1.228-0.0591{\text{ pH }}+{\frac {0.0591}{4}}\log {P_{\mathrm {O_{2}} }}}

giving:

2 N2 + 5 O2 + 2 H2O → 4 NO3− + 4 H+ 0 = 0.018 − 0.0118 pH + 0.0591 10 log ⁡ ( N O 3 − ) 2 P N 2 − 0.0591 4 log ⁡ P O 2 {\displaystyle \qquad 0=0.018-0.0118{\text{ pH }}+{\frac {0.0591}{10}}\log {\frac {(\mathrm {NO_{3}^{-}} )^{2}}{P_{\mathrm {N_{2}} }}}-{\frac {0.0591}{4}}\log {P_{\mathrm {O_{2}} }}}

Dividing by 0.0118 and rearranging gives the equilibrium relation:

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