Sulphides in Wastewater: Why the Chemistry Behind H2S Matters for Treatment Decisions

Sulphides are compounds of sulphur in the −2 oxidation state, and metal sulphides can be regarded as salts of hydrosulphuric acid. In a wastewater context that definition matters less than what follows from it: sulphides are reducing, they hydrolyse to an alkaline solution, they release a toxic gas when acidified, and they precipitate a long list of metals. Almost every operational problem associated with them traces back to one of those four behaviours.

How they form and how they are made

Metal sulphides can be prepared by reacting a metal directly with sulphur, by passing hydrogen sulphide gas through a metal-salt solution, or by adding sodium sulphide to a salt solution. Industrially, inorganic sulphides come from a handful of routes. Direct combination of the elements gives carbon disulphide, written C + 2S → CS2. Reduction of sulphates or higher sulphides with carbon yields sodium sulphide. Metathesis reactions in solution or at high temperature, preparation from thio-salts, and thermal decomposition of higher-valence sulphides round out the list; molybdenum trisulphide decomposes on heating to the disulphide and elemental sulphur. Ammonium salts such as NH4Cl and (NH4)2MoO4 are common starting materials in the thio-salt routes.

Solubility, and why pH drives everything

Alkali-metal sulphides and ammonium sulphide dissolve readily in water and give alkaline solutions through hydrolysis. Alkaline-earth sulphides, and those of scandium, yttrium and the lanthanides, are relatively insoluble. Where the cation carries an 18 or 18+2 outer-electron configuration, strong polarization tends to produce insoluble, strongly coloured sulphides.

Most water-insoluble sulphides dissolve in acid with evolution of hydrogen sulphide; a few of the most insoluble, CuS and HgS among them, need an oxidising acid, which oxidises the sulphur to elemental sulphur. Insoluble sulphides sit in a dissolution–precipitation equilibrium, and by controlling acidity the concentration of S²⁻ can be moved at will. That is the entire basis of the classical hydrogen-sulphide scheme for separating and identifying metal ions, and it is why a shift in pH in a collection system can strip metals out of solution or put them back.

Hydrolysis and alkalinity

Sulphide ions hydrolyse in two steps, written S2- + H2O ↔ HS- + OH- and then HS- + H2O ↔ H2S + OH-. With H₂S pKa values of about 6.89 and 15±2, sulphide solutions show alkalinity that in the alkali-metal case can rival the corresponding hydroxide. Any plant measuring alkalinity upstream of a sulphide source should expect that contribution rather than chasing an instrument fault.

Redox behaviour

Sulphur in the −2 state is reducing, and depending on conditions oxidises to elemental sulphur, sulphite or sulphate. The half reaction S2- - 2e- = S is quoted at −0.407V, and 0.407V is the value analysts cite for the couple. In practice this means sulphides scavenge oxidant: any oxidation step, whether chlorination, peroxide or ozone, will consume part of its dose on sulphide before it touches the target pollutant.

Roasting sulphide minerals follows one of two paths: conversion to the oxide with sulphur released as sulphur dioxide, as in the galena-to-lead step 2PbS + 3O₂ → 2PbO + 2SO₂, or oxidation to a soluble sulphate. Both are industrial routes to winning metal from sulphide ore, and both are reminders of how much oxidant a sulphide load represents.

Hydrogen sulphide: the exposure and detection problem

H₂S is a colourless, toxic gas with the familiar rotten-egg smell, and its permissible concentration in air is no more than 0.01 mg/L. It binds the ferrous ions in human haemoglobin to form ferrous sulphide and renders them inactive; repeated exposure produces chronic effects including dulled smell, weight loss and headaches. In the laboratory it is usually generated by reacting a metal sulphide with acid. Its aqueous solution is hydrosulphuric acid, a diprotic weak acid, and it is strongly reducing whether the medium is acidic or alkaline.

Two spot tests are sensitive enough for routine checks. Mixing an alkaline solution of a soluble sulphide with 1% sodium nitroprusside produces a red-purple colour of varying depth when S²⁻ is present, with a sensitivity around 1:50,000; the colour comes from formation of the [Fe(CN)₅(NOS)]⁴⁻ ion. Alternatively, adding concentrated HCl, a few crystals of p-aminodimethylaniline and 0.1 mol/L ferric chloride to the sample gives a blue colour after 2–3 minutes through formation of methylene blue.

Because H₂S gas is both highly toxic and awkward to store, thioacetamide, CH₃CSNH₂ or TAA, is generally used as the precipitant instead. Its hydrolysis depends on the medium: in acid it yields H₂S, in ammoniacal solution it yields hydrosulphide in place of ammonium sulphide, and in alkaline solution it yields sulphide in place of sodium sulphide. The equations are CH₃CSNH₂ + H⁺ + 2H2O ↔ CH₃COOH + NH₄⁺ + H₂S in acid, CH₃CSNH₂ + 2NH3 ↔ CH₃–C(=NH)NH₂ + NH₄⁺ + HS⁻ in ammoniacal solution, and CH₃CSNH₂ + 3OH⁻ ↔ CH₃COO⁻ + NH3 + H2O + S in alkali.

Polysulphides and what acidification does

Polysulphides contain ions Sₙ with n running from 2 to 9, made by boiling sulphur in a sulphide solution. Their solutions are yellow, deepening as n increases. The ions carry persulphide bonds similar to peroxides — oxidising, though less so than peroxide — and they disproportionate. They are very unstable in acid: acidification releases hydrogen sulphide and elemental sulphur, written simply as Sn2- + 2H → H2S + (n-1)S. That is the reaction that turns an innocuous-looking alkaline stream into a gas hazard the moment it is mixed with an acidic one.

Polysulphide ions also act as ligands; sodium polysulphide reacting with a titanocene dichloride forms a coordination compound containing a TiS₅ ring.

Where sulphides show up industrially

Sulphide minerals number around 350 species, more than two-thirds of them true sulphides, the rest selenides, tellurides, arsenides and a few antimonides and bismuthides. The class is only 0.15% of crustal mass, most of it iron sulphides, but it concentrates into ores that supply most of the world's copper, lead, zinc, mercury, antimony, bismuth, molybdenum, nickel and cobalt.

Beyond metallurgy the applications are broad. Molybdenum disulphide catalyses hydrogenation of sulphur-containing organics where ordinary catalysts would be poisoned. Cadmium sulphide serves in photoelectric cells and, at high purity, as a semiconductor; zinc and cadmium sulphides make phosphors. Lead sulphide goes into infrared sensors. Calcium, barium and ammonium polysulphides are fungicides and insecticides. Carbon disulphide is an industrial solvent, a route to carbon tetrachloride, and a reagent for inserting the –C(=S)–S– group. Sodium sulphide is consumed in sulphur dyes, organic drugs and pulp; calcium and barium sulphides in luminous paints; tetraphosphorus trisulphide in matches and fireworks.