Iron(II) sulfate, anhydrous

Every photographic formula that says “ferrous sulphate” means the green crystals, not this. The anhydrous salt is the same iron(II) with the water driven out of the lattice, and the difference is not academic: 151.91 against 278.02 means that a spoonful of the dry white powder carries 1.83 times as much iron as the same weight of crystals, on the course’s own arithmetic from the two molar masses. That ratio is the whole reason this page exists separately.

In photography

Iron(II) is a one-electron reducing agent, and that is its entire photographic career. Every use below is the same reaction seen from a different angle: Fe²⁺ gives up one electron, becomes Fe³⁺, and something else is reduced.

⚗The chemistry

In wet collodion the something else is dissolved silver. Hardwich’s 1864 manual sets out the principle before writing any equation: developers “have a tendency, when mixed with soluble Salts of Silver, to set Silver free”, and salts of iron do it by uniting with the acid and the oxygen of the silver salt. Because it takes one electron to turn Ag⁺ into an atom of silver, and iron(II) supplies exactly one, the arithmetic is one for one.

Ag⁺ + Fe²⁺ → Ag + Fe³⁺

The one-electron step underneath every iron developer

Where the silver comes from decides what kind of development it is. On a collodion plate the free silver nitrate clinging to the wet film is the reservoir, so the image is built by silver arriving from solution and settling on the light-struck specks — physical development, not the chemical development of a gelatin emulsion in which the developer attacks the halide grain itself. The distinction is why an iron developer is poured on a plate that has not been rinsed, and why letting the plate dry ends the process.

Making ferrous oxalate. The second use of the salt is as a source of iron(II) for a developer that is not an iron sulfate developer at all. Wall’s 1924 formulary precipitates ferrous oxalate by mixing 278 parts of ferrous sulfate in 500 parts of water with 184 parts of neutral potassium oxalate in 500 parts of water, washing the precipitate with four or five changes of distilled water and drying it: “The result should be 180 parts of ferrous oxalate.” Those 278 parts are the heptahydrate’s formula weight, and the anhydrous salt at 151.91 would have to be weighed at 0.546 times that figure — 152 parts — to supply the same iron.

What more or less of it does. In the collodion developer, more iron means a faster, denser, flatter build-up and a greater risk of fog; the historical control is not less iron but more acid. Towler’s 1864 formulas make the point by contrast: three drachms of the crystals in four ounces of water with three drachms of acetic acid for an ambrotype, two drachms in thirty-two drachms of water for a negative, and an explicit instruction that in warm weather or after a long exposure more acid is needed “to keep the reducing agent in check”, while for very short exposures the iron may be used with no acid at all.

Properties

The ChEBI description PubChem carries sets out the hydrate ladder that governs everything practical about this substance: the heptahydrate is the form that occurs naturally, and it loses water to give the tetrahydrate at 57 °C and the monohydrate at 65 °C. CAMEO’s datasheet for ferrous sulfate describes a greenish or yellow-brown crystalline solid melting at 64 °C and losing its seven waters at 90 °C. HSDB records that the monohydrate is a white-to-yellow crystalline powder that loses water at about 300 °C and decomposes above that.

Which powder is in the jar. The anhydrous salt is a white hygroscopic solid; the photographic salt is blue-green. Colour is therefore a usable check, and so is weight: a tub of anhydrous iron(II) sulfate left open will gain mass as it hydrates, and a tub of the heptahydrate left in warm dry air will lose it. Neither is what the balance assumed.

Air is the other problem. Hardwich records that pure crystals “gradually absorb Oxygen and become rusty on the surface”, that a solution colourless at first turns red and deposits a brown powder — a basic iron(III) sulfate — and that adding sulfuric or acetic acid prevents the deposit because the brown powder is soluble in acid. Wall’s 1912 dictionary gives the same instruction in one line: crystals covered with brownish rusty powder have already oxidised, and a solution whose colour has changed should be rejected. Wall’s 1924 stock solution carries 1 cm³ of sulfuric acid per litre and the instruction to keep it in small bottles filled to the cork.

Handling

The aggregated ECHA notifications held by PubChem classify iron(II) sulfate Warning, irritant: harmful if swallowed, skin irritation and serious eye irritation, each in 99.8 per cent of the 2,126 reports that carry hazard codes. CAMEO’s health hazard entry describes the consequence of swallowing it — abdominal pain, retching, diarrhoea, dehydration, shock, pallor, cyanosis, a rapid or weak pulse. That is a description of iron poisoning, and it is the reason this substance is treated as something to keep out of the mouth rather than as an inert salt.

Why Level A. Against the course rubric this substance meets the Level A criteria and no more. Its classification tops out at irritant and harmful if swallowed, which is exactly the ceiling Level A sets; nothing in the way the course handles it is heated above 50 °C; and the waste is spent developer that can be collected. NIOSH sets a recommended limit of 1 mg/m³ for soluble iron salts measured as iron, and HSE’s EH40 gives the identical British long-term figure with 2 mg/m³ over fifteen minutes, so weighing the dry powder is done without raising dust — a control, not a prohibition. Two Level B habits are added by preference: splash goggles rather than glasses for the dry salt, because the eye statement is near unanimous, and hand washing before eating, because ingestion is the route that matters. The procedures the salt appears in may themselves sit higher: wet collodion is governed by its silver nitrate, its ether and its cadmium bromide, not by the iron.

⚠Safety

Weigh over a tray with the container away from the face; powder into water, never water into powder. Gloves and eye protection whenever a solution is handled. CAMEO records iron(II) sulfate as a weak inorganic reducing agent that reacts with oxidising agents to generate heat and products that may be flammable or otherwise reactive, so it is kept away from bleaches, peroxides, persulfates and chlorates. Acidify stock solutions and keep them full and stoppered: the brown deposit that forms otherwise is basic iron(III) sulfate, and it stains gelatin and paper.

Waste

An iron developer leaves the darkroom as a dilute solution of iron(II) and iron(III) sulfate with whatever acid was in it, and, from a collodion plate, dissolved silver picked up from the film. The silver decides the route: it is recovered from the first baths before anything else is considered, and the course treats every silver-bearing bath as a recovery stream rather than a disposal problem. What remains is iron. CAMEO’s datasheet is blunt that the primary hazard of ferrous sulfate is the threat to the environment, and Japan’s NITE classification of the heptahydrate carries H402, harmful to aquatic life, so a spent bath does not go onto a garden, into a soakaway or into a watercourse. Kodak’s J-300 guidance for amateur photographers sends developers, stop baths, fixers after silver recovery and wash water to the sewer where a household has one, and is explicit that septic systems are not designed to treat photographic processing solutions. ILFORD tells domestic users in the United Kingdom to bottle wastes separately, label them and take them to a household waste and recycling centre’s chemical cupboard. Check your local regulations; they govern, and they differ.

History

Hardwich’s 1864 manual places the salt in its industrial context before its photographic one: it is “an abundant substance, and is used for a variety of purposes in the arts”, prepared on such a scale that commercial material “requires recrystallization to render it sufficiently pure for Photographic purposes”. Its photographic career began in the 1840s. Hunt’s Researches on Light quotes a contemporary account of the collodion process crediting “the photographic property of iodide, which was discovered by Dr. Woods of Parsonstown, in Ireland, and that of sulphate of iron, for which science is indebted to the researches of Mr. Robert Hunt” — Hunt’s own claim, reported at second hand in his own book, which is how much of the period’s attribution works. Archer’s 1854 manual of the collodion process gives the salt its period identity in three sentences: obtained by the action of dilute sulfuric acid on metallic iron, decomposing in solution unless a minute portion of sulfuric acid is present, and “a cheap and energetic developing agent”.

What replaced it was not a better iron salt but a different chemistry. Wall’s 1924 verdict on the ferrous oxalate developer that iron(II) sulfate feeds is that its use “has been completely abandoned”, because plates are actually slower in it than in the organic developers and because calcium oxalate and basic iron salts deposit in the gelatin. The iron developer survives in exactly the place where its physical development is the point rather than the problem: the wet collodion plate, where there is free silver nitrate on the film and no gelatin to stain.