PbII

PbII, written Pb2+ in chemical notation, denotes lead in the divalent oxidation state. It is the most common oxidation state of lead in many inorganic and coordination chemistry contexts, especially under reducing or neutral conditions. In aqueous solution, Pb2+ is typically solvated by water molecules, but its lone pair on the lead center can produce stereochemically active distortions in coordination geometry, leading to structures that deviate from ideal octahedral symmetry. In chloride-rich media, Pb2+ readily forms complex ions such as [PbCl3]- and [PbCl4]2-, which influence solubility and reactivity.

Pb(II) forms a wide range of salts. Salts with halides (for example PbCl2, PbBr2) often have limited solubility in water, though solubility can increase with temperature and complexation. Lead(II) also undergoes hydrolysis in water, generating various Pb(OH)x species, and at neutral to basic pH Pb(OH)2 tends to precipitate. The chemistry of Pb2+ is further influenced by the presence of other ligands, including sulfates, carbonates, and organic molecules, which can stabilize different species through complexation.

Redox chemistry of lead commonly involves interconversion between Pb2+ and metallic lead (Pb0) or, less stably, Pb4+. Pb(IV) compounds are typically strong oxidizers and less stable in aqueous solutions, with Pb(IV) reduction occurring under milder conditions in some environments.

Occurrence and uses: major lead minerals include cerussite (PbCO3) and anglesite (PbSO4). Historically, Pb(II) compounds were used as pigments (white lead) and in various commercial applications; today, Pb2+ remains central to lead-acid battery technology and to certain catalysts and materials. Toxicology: Pb2+ is highly toxic, affecting nervous, renal, and hematological systems, and exposure is tightly regulated worldwide.