Nitrite

The ** nitrite ** [|ion] has the [|chemical formula] NO2−. The [|anion] is symmetric with equal N-O bond lengths and a O-N-O bond angle of about 120°. On protonation the unstable weak acid [|nitrous acid] is produced. Nitrite can be oxidized or reduced, with product somewhat dependent on the oxidizing/reducing agent. The nitrite ion is an [|ambidentate ligand] and is known to bond to metal centers in at least five different ways. [|[1]] Nitrite is important in biochemistry as a source of the [|vasodilator] [|nitric oxide]. Nitrites are used for curing [|meat]. In organic chemistry the NO2 group is present in nitrous acid esters and nitro compounds.



Nitrite in biochemistry
Sodium nitrite is used for the [|curing] of [|meat] because it prevents bacterial growth and, in a reaction with the meat's [|myoglobin], gives the product a desirable dark red color. Because of the toxicity of nitrite (the lethal dose of nitrite for humans is about 22 mg per kg body weight), the maximum allowed nitrite concentration in meat products is 200 [|ppm]. Under certain conditions, especially during cooking, nitrites in meat can react with degradation products of [|amino acids], forming [|nitrosamines] , which are known [|carcinogens]. [|[4]]

Nitrite is detected and analyzed by the [|Griess Reaction], involving the formation of a deep red-colored [|azo dye] upon treatment of a NO2−-containing sample with [|sulfanilic acid] and naphthyl-1-amine in the presence of acid. [|[5]] Nitrite can be reduced to [|nitric oxide] or [|ammonia] by many species of [|bacteria]. Under hypoxic conditions, nitrite may release nitric oxide, which causes potent [|vasodilation]. Several mechanisms for nitrite conversion to NO have been described including enzymatic reduction by [|xanthine oxidoreductase], [|nitrite reductase] and [|NO synthase] (NOS), as well as nonenzymatic acidic [|disproportionation].

**nitrite****,** any member of either of two classes of compounds derived from [|nitrous acid]. Salts of nitrous acid are [|ionic compounds] containing the nitrite ion, NO -2, and a [|positive ion] such as Na + in sodium nitrite (NaNO 2 ). Esters of nitrous acid are [|covalent compounds] having the structure R−O−N−O, in which R represents a carbon-containing combining group such as ethyl (C 2 H 5 ) in ethyl nitrite. These covalent nitrites are isomers of the [|nitro compounds] —//e.g.,// nitroethane—which are considered to be derivatives of [|nitric acid] rather than of nitrous acid. Nitrites usually are prepared by absorption of [|nitric oxide] and [|nitrogen dioxide] in an [|alkaline solution]. In an older method, [|sodium nitrate] was fused with lead, and the resulting sodium nitrite was dissolved in water and separated from the by-product, lead oxide, by filtration. Nitrites are used as food preservatives and in medicine as vasodilators to relieve cardiac pain. //See also// [|nitroso compound].



Nitrite salts
[|Sodium nitrite] is made industrially by passing //nitrous fumes// into aqueous [|sodium hydroxide] or [|sodium carbonate] solution. [|[1]] NO + NO2 + 2NaOH (or Na2CO3) → 2NaNO2 +H2O ( or CO2) The product is purified by recrystallization. Alkali metal nitrites are thermally stable up to and beyond the melting point (441 °C for KNO2). [|Ammonium nitrite] can be made from [|dinitrogen trioxide], N2O3, which is formally the [|anhydride] of nitrous acid. 2NH3 + H2O +N2O3 → 2NH4NO2 This compound may decompose explosively on heating. In organic chemistry nitrites are used in [|diazotization] reactions.

Structure
The two [|canonical structures] of NO2−, which contribute to the resonance hybrid for the nitrite ion The nitrite ion has a symmetrical structure (C2v [|symmetry] ) with both N-O bonds having equal length. In [|valence bond theory] it is described as a [|resonance hybrid] with equal contributions from two canonical forms that are mirror images of each other. In [|molecular orbital theory] there is a [|sigma bond] between each oxygen atom and the nitrogen atom, and a delocalized [|pi bond] made from the [|p orbitals] on nitrogen and oxygen atoms which are perpendicular to the plane of the molecule. The negative charge of the ion is equally distributed on the two oxygen atoms. Both nitrogen and oxygen atoms carry a [|lone pair] of electrons. Therefore the nitrite ion is a [|Lewis base]. Moreover, it can act as an [|ambidentate ligand] towards a metal ion, donating a pair of electrons from either nitrogen or oxygen atoms.

Acid-base properties
Dimensions of //trans// HONO (from the [|microwave spectrum] ) In aqueous solution [|nitrous acid] is a [|weak acid]. HNO2 H+ + NO2-; [|pKa] = ca. 3.3 at 18 °C [|[2]] Nitrous acid is volatile; in the gas phase it exists predominantly as a //trans//- planar molecule. In solution it is unstable with respect to the [|disproportionation] reaction 3HNO2 (aq) H3O+ + NO3- + 2NO This reaction is slow at 0 °C. [|[1]] 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.

Oxidation and reduction
The formal [|oxidation state] of the nitrogen atom in a nitrite is +3. This means that it is can be either oxidised 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. [|[3]] The data can be extended to include products in lower oxidation states. For example, H2N2O2 + 2H+ + 2//e//- N2 + 2H2O; E0 = 2.65V Oxidation reactions usually result in the formation of the [|nitrate] ion, with nitrogen in oxidation state +5. For example, oxidation with [|permanganate] can be used for quantitative analysis of nitrite, by titration. 5NO2- + 2MnO4- + 6H+ → 5NO3- + 2Mn2+ + 3H2O The product of reduction reactions are various depending on the reducing agent used. 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 occurs with [|hydrogen sulfide]. With the [|hydrazinium] cation, N2H5+, [|hydrogen azide], HN3, is produced HNO2 + N2H5+ → 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. [|[1]]
 * ~ Half-reaction ||~ E0/V ||
 * NO3- + 3H+ + 2//e//- [[image:http://upload.wikimedia.org/wikipedia/commons/thumb/9/96/Equilibrium.svg/15px-Equilibrium.svg.png width="15" height="13" caption="is in equilibrium with"]] HNO2 + H2O || +0.94 ||
 * 2HNO2+ 4H+ + 4//e//- [[image:http://upload.wikimedia.org/wikipedia/commons/thumb/9/96/Equilibrium.svg/15px-Equilibrium.svg.png width="15" height="13" caption="is in equilibrium with"]] H2N2O2 + 2H2O || +0.86 ||
 * N2O4 + 2H+ + 2//e//- [[image:http://upload.wikimedia.org/wikipedia/commons/thumb/9/96/Equilibrium.svg/15px-Equilibrium.svg.png width="15" height="13" caption="is in equilibrium with"]] 2HNO2 || +1.065 ||
 * 2HNO2+ 4H+ + 4//e//- [[image:http://upload.wikimedia.org/wikipedia/commons/thumb/9/96/Equilibrium.svg/15px-Equilibrium.svg.png width="15" height="13" caption="is in equilibrium with"]] N2O + 3H2O || +1.29 ||