Nitrits

Nitrites, or nitrite ions, refer to the chemical species NO2−, the conjugate base of nitrous acid (HNO2). They occur as salts, such as sodium nitrite (NaNO2) and potassium nitrite (KNO2), and as esters in organic chemistry. In aqueous solution nitrite exists in equilibrium with nitrous acid, with a pKa around 3.3, so the proportion of protonated species depends on pH.

Nitrites are common in the environment as part of the nitrogen cycle. They are mainly formed by microbial reduction of nitrate and can be further oxidized back to nitrate. In water and soil, nitrite levels can vary with microbial activity, oxygen status, and agricultural inputs. In humans, nitrite can be produced endogenously under certain conditions and can act as a reservoir for nitric oxide, especially in hypoxic tissues.

Chemically, nitrite is a weak base and can function as both an oxidizing and a reducing agent, depending on context. It participates in nitrosation reactions, forming N-nitroso compounds when it encounters secondary amines in acidic environments. This reactivity is relevant to food chemistry and public health, as it can lead to the formation of nitrosamines under certain processing conditions.

Uses and regulation: Nitrite salts are widely used in food processing to preserve and color cured meats and to inhibit bacterial growth, notably Clostridium botulinum. They also serve as reagents in organic synthesis and have industrial applications. Regulatory frameworks restrict nitrite levels in drinking water and foods to minimize health risks. Typical limits in water are on the order of 1 mg/L as nitrite, with additional controls in dietary products.

Health considerations: In humans, nitrite exposure can cause methemoglobinemia, particularly in infants consuming high-nitrite water or foods. There is concern about the potential formation of carcinogenic N-nitrosamines under certain conditions, leading to processing controls and monitoring to mitigate risks.