arsenate

Arsenate refers to the chemical species and salts derived from arsenic acid (H3AsO4) that contain the arsenate anion AsO4^3-. In its most common oxidation state, arsenic forms a tetrahedral AsO4^3- unit and acts as a chemical analog of phosphate (PO4^3-). Arsenate salts occur with a variety of cations, for example sodium arsenate Na3AsO4 and calcium arsenate Ca3(AsO4)2, and are encountered in minerals, soils, and water.

In aqueous solution, arsenate exists mainly as H2AsO4^- and HAsO4^2-, with the relative distribution governed by pH. The acid–base steps H3AsO4 ⇌ H2AsO4^- + H^+ (pKa1 ≈ 2.2), H2AsO4^- ⇌ HAsO4^2- + H^+ (pKa2 ≈ 6.9), and HAsO4^2- ⇌ AsO4^3- + H^+ (pKa3 ≈ 11.5).

Geochemically, arsenate occurs in arsenate minerals and enters the environment via natural weathering and from human activities such as mining and ore processing. Groundwater contamination by arsenate arises from oxidation of arsenite-bearing waters or dissolution from arsenic-bearing minerals; drinking water standards in many countries regulate total arsenic, including arsenate, to protect health.

Biologically, arsenate behaves as a phosphate analog and can interfere with energy metabolism. In some organisms, arsenate can replace phosphate in biochemical intermediates to form unstable arsenate esters that readily hydrolyze, yielding little or no net ATP and disrupting normal cellular processes. Historically, arsenate compounds were used as pesticides and pigments, though their use has declined due to toxicity concerns.

Health and environmental concerns drive treatment strategies for arsenate-contaminated water, including adsorption onto iron oxides, coagulation and filtration, ion exchange, and reverse osmosis. Regulatory limits for arsenic in drinking water are commonly around 10 micrograms per liter, with stricter standards in some regions.