Iron(II) sulfate

Iron(II) sulfate heptahydrate prepared from steel and sulfuric acid (looks more blue in person)

Iron(II) sulfate, also known as ferrous sulfate is the sulfate salt of the iron(II) ion. It is most commonly seen as the heptahydrate which forms blue-green crystals. It is a brownish-white powder when anhydrous.



Iron(II) sulfate is an easy source of iron(II) ions in solution, since it is readily available and not expensive.

Upon standing in air, iron(II) sulfate will oxidize to a mixture of iron(III) sulfate and iron(III) oxide because iron(II) compounds are not stable when not kept at a low pH. This can be prevented by adding a small amount of sulfuric acid. When heated to 680°C, iron(II) sulfate begins to decompose, releasing sulfur dioxide and sulfur trioxide, leaving behind iron(III) oxide. It also reacts with hydrogen peroxide.

2 FeSO4 → Fe2O3 + SO2 + SO3


Iron(II) sulfate is usually seen as the heptahydrate, which forms blue-green crystals. When heated to around 300°C, it loses all of its water of crystallization[1]


Iron(II) sulfate heptahydrate can be found at some garden stores, and can also be bought cheaply online.[2] Its purity when bought can be told by it's color, impure samples having a dark green color and brown to gray hue.


Iron(II) sulfate can be prepared with iron or steel scraps and dilute sulfuric acid. Concentrated sulfuric acid will not work. If steel is used, the carbon must be filtered out after the reaction is complete. Do not leave the solution to crystallize by evaporation, as the product will become oxidized and impure. Instead, it must be heated without boiling until crystals are visible, cooled, and then dried in a dessicator.[3]

Iron(II) sulfate is produced by addition of iron to copper(II) sulfate. During this reaction, the carbon from steel will leach, forming a black goo between the copper layer and iron. Copper(II) oxide will also form, which, because it's also black, will make it difficult to determine how much metallic copper was oxidized. Eliminating the air from water prior to the reaction or adding a very small quantity of acid will reduce the formation of the copper oxide and increase the yield.




Wet iron sulfate should not be handled directly, as it may contain excess sulfuric acid that can burn the skin, if an excess of acid was used. This is not an issue if the sulfate was prepared with copper(II) sulfate and iron metal.


Ferrous sulfate should be stored in closed bottles, away from moisture.


Ferrous sulfate does not require special disposal.



Relevant Sciencemadness threadsEdit

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