Nitrite

Nitrite is the conjugate base of the weak acid nitrous acid:

HNO₂ ⇌ H⁺ + NO⁻
₂;      pK_a ≈ 3.3 at 18 °C[3]

Nitrous acid is also highly volatile, tending to disproportionate:

3 HNO₂ (aq) ⇌ H₃O⁺ + NO⁻
₃ + 2 NO

This reaction is slow at 0 °C.[2] Addition of acid to a solution of a nitrite in the presence of a reducing agent, such as iron(II), is a way to make nitric oxide (NO) in the laboratory.

The formal oxidation state of the nitrogen atom in nitrite is +3. This means that it can be either oxidized to oxidation states +4 and +5, or reduced to oxidation states as low as −3. Standard reduction potentials for reactions directly involving nitrous acid are shown in the table below:[4]

Half-reaction	E0 (V)
NO− 3 + 3 H+ + 2 e− ⇌ HNO2 + H2O	+0.94
2 HNO2 + 4 H+ + 4 e− ⇌ H2N2O2 + 2 H2O	+0.86
N2O4 + 2 H+ + 2 e− ⇌ 2 HNO2	+1.065
2 HNO2+ 4 H+ + 4 e− ⇌ N2O + 3 H2O	+1.29

The data can be extended to include products in lower oxidation states. For example:

H₂N₂O₂ + 2 H⁺ + 2 e⁻ ⇌ N₂ + 2 H₂O;      E⁰ = +2.65 V

Oxidation reactions usually result in the formation of the nitrate ion, with nitrogen in oxidation state +5. For example, oxidation with permanganate ion can be used for quantitative analysis of nitrite (by titration):

5 NO⁻
₂ + 2 MnO⁻
₄ + 6 H⁺ → 5 NO⁻
₃ + 2 Mn²⁺ + 3 H₂O

The product of reduction reactions with nitrite ion are varied, depending on the reducing agent used and its strength. With sulfur dioxide, the products are NO and N₂O; with tin(II) (Sn²⁺) the product is hyponitrous acid (H₂N₂O₂); reduction all the way to ammonia (NH₃) occurs with hydrogen sulfide. With the hydrazinium cation (N
₂H⁺
₅) the product of nitrite reduction is hydrazoic acid (HN₃), an instable and explosive compound:

HNO₂ + N
₂H⁺
₅ → HN₃ + H₂O + H₃O⁺

which can also further react with nitrite:

HNO₂ + HN₃ → N₂O + N₂ + H₂O

This reaction is unusual in that it involves compounds with nitrogen in four different oxidation states.[2]

Nitrite is detected and analyzed by the Griess Reaction, involving the formation of a deep red-colored azo dye upon treatment of a NO⁻
₂-containing sample with sulfanilic acid and naphthyl-1-amine in the presence of acid.[5]

Nitrite is an ambidentate ligand and can form a wide variety of coordination complexes by binding to metal ions in several ways.[2] Two examples are the red nitrito complex [Co(NH₃)₅(ONO)]²⁺ is metastable, isomerizing to the yellow nitro complex [Co(NH₃)₅(NO₂)]²⁺. Nitrite is processed by several enzymes, all of which utilize coordination complexes.

Azo dyes and other colorants are prepared by the process called diazotization, which requires nitrite.[1]

Sodium nitrite is used to speed up the curing of meat and also impart an attractive colour.[8] A 2018 study by the British Meat Producers Association determined that legally permitted levels of nitrite have no effect on the growth of the Clostridium botulinum bacteria which causes botulism, in line with the UK's Advisory Committee on the Microbiological Safety of Food opinion that nitrites are not required to prevent Clostridium botulinum growth and extend shelf life.[9] In the U.S., meat cannot be labeled as "cured" without the addition of nitrite.[10][11][12] In some countries, cured-meat products are manufactured without nitrate or nitrite, and without nitrite from vegetable source. Parma ham, produced without nitrite since 1993, was reported in 2018 to have caused no cases of botulism.[8]

In mice, food rich in nitrites together with unsaturated fats can prevent hypertension, which is one explanation for the apparent health effect of the Mediterranean diet.[13] However, adding nitrites to meat has been shown to generate known carcinogens; the World Health Organization (WHO) advises that eating 50 g (1.8 oz) of nitrite processed meat a day would raise the risk of getting bowel cancer by 18% over a lifetime.[8] The recommended maximum limits by the World Health Organization in drinking water are 3 mg L⁻¹ and 50 mg L⁻¹ for nitrite and nitrate ions, respectively.[14]

In a reaction with the meat's myoglobin, nitrite gives the product a desirable pink-red "fresh" color, such as with corned beef. In the US, nitrite has been formally used since 1925. According to scientists working for the industry group American Meat Institute, this use of nitrite started in the Middle Ages.[15] Historians[16] and epidemiologists[17] argue that the widespread use of nitrite in meat-curing is closely linked to the development of industrial meat-processing. French investigative journalist Guillaume Coudray asserts that the meat industry chooses to cure its meats with nitrite even though it is established that this chemical gives rise to cancer-causing nitroso-compounds.[18] Nitrite is commonly added to fermented meat to speed up the curing and also impart an attractive colour while having no effect on the growth of the Clostridium botulinum bacteria which causes botulism.[19][20] Some traditional and artisanal producers avoid nitrites.

Nitrites in the form of sodium nitrite and amyl nitrite are components of many cyanide antidote kits.[21] Both of these compounds bind to hemoglobin and oxidize the Fe²⁺ ions to Fe³⁺ ions forming methemoglobin. Methemoglobin in turn binds to cyanide (CN) creating cyanmethemoglobin, effectively removing cyanide from the complex IV of the electron transport chain (ETC) in mitochondria, which is the primary site of disruption caused by cyanide. Another mechanism by which nitrites are useful in treating cyanide toxicity is the generation of nitric oxide (NO). NO displaces the CN from the cytochrome c oxidase (ETC complex IV), making it available for methemoglobin to bind.[22]