Nitrate
Chemical and biochemical properties

In the NO − 3 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 (NH NO), or black powder, ignited by the shock wave of a primary explosive. However, in contrast to red fuming nitric acid (HNO /N O), or concentrated nitric acid (HNO), nitrate dissolved in aqueous solution at neutral or high pH is only a weak oxidizing agent and is stable under sterile, or aseptic, conditions, in the absence of microorganisms. To increase its oxidation power, acidic conditions and high concentrations are needed, under which nitrate transforms into nitric acid. This behaviour is consistent with the general theory of reduction-oxidation (redox) in electrochemistry: oxidizing power is exacerbated under acidic conditions while the power of reducing agents is reinforced under basic conditions. This can be illustrated by means of a Pourbaix diagram (E –pH diagram) drawn using the Nernst equation and the corresponding redox reactions. During the reduction of oxidizers, the oxidation state decreases and oxide ions (O) in excess released in water by the reaction are more easily protonated under acid conditions (O + 2 H → H O) which drives the reduction reaction to the right according to Le Chatelier's principle. For the oxidation of reducing agents, the reverse occurs: as the oxidation state increases, oxide anions are needed to neutralise the surplus of positive charges born by the central atom. As basic conditions favor the production of oxide anions (2 OH → O + H O), this drives the chemical equilibrium of the oxidation reaction to the right.

Meanwhile, nitrate is used as a powerful terminal electron acceptor by denitrifying bacteria to deliver the energy they need to thrive. Under anaerobic conditions, nitrate is the strongest electron acceptor used by prokaryote microorganisms (bacteria and archaea) to respirate. The redox couple NO − 3 / N is at the top of the redox scale for the anaerobic respiration, just below the couple oxygen (O / H O), but above the couples Mn(IV)/Mn(II), Fe(III)/Fe(II), SO 2− 4 / HS, CO / CH. In natural waters, inevitably contaminated by microorganisms, nitrate is a quite unstable and labile dissolved chemical species because it is metabolised by denitrifying bacteria. Water samples for nitrate/nitrite analyses need to be kept at 4 °C in a refrigerated room and analysed as quick as possible to limit the loss of nitrate.

In the first step of the denitrification process, dissolved nitrate (NO − 3) is catalytically reduced into nitrite (NO − 2) by the enzymatic activity of bacteria. In aqueous solution, dissolved nitrite, N(III), is a more powerful oxidizer that nitrate, N(V), because it has to accept less electrons and its reduction is less kinetically hindered than that of nitrate.

During the biological denitrification process, further nitrite reduction also gives rise to another powerful oxidizing agent: nitric oxide (NO). NO can fix on myoglobin accentuating its red coloration. NO is an important biological signaling molecule and intervenes in the vasodilation process, but it can also produce free radicals in biological tissues, accelerating their degradation and aging process. The reactive oxygen species (ROS) generated by NO contribute to the oxidative stress, a condition involved in vascular dysfunction and atherogenesis.
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