Pure iron is a silvery gray/blue metal, which is oxidized easily. It is slowly attacked by water, especially when the water contains some dissolved electrolyte (e.g. salt). This attack of iron by water is the well-known rusting of iron. By dilute acids, it is attacked more quickly.
The metal can be purchased in powder form from several vendors on eBay. This powdered metal can be interesting for many kinds of experiments with magnets, and it can be used as a source for creation of solutions of iron-salts. For the home chemist, however, it is more convenient to have a water-soluble compound of this metal at hand.
The most stable oxidation states of iron in its compounds are +2 and +3. Iron also can be brought to the +6 oxidation state in the so called ferrate ion and probably it exists in the +4 oxidation state as well in a transient state, on addition of hydrogen peroxide to a solution of an iron (II) salt.
The general public has access to quite a large number of iron compounds. This is due to the fact that iron salts are used for many processes in photography, but also in other applications, such as ink-making, and old processes, such as making cyanotypes. For the home chemist it is not necessary to have all these iron compounds at hand. Having one soluble iron (II) compound and one iron (III) compound, plus the hexacyanoferrates is sufficient.
Below follows a list of available iron-compounds:
Ferrous sulfate is available in many grades. The commercially available compound usually is a greenish crystalline solid with a yellow/brown tinge. This is fairly impure ferrous sulfate, which also contains quite some basic ferric sulfate. For certain experiments with iron (II) this is OK, but frequently, this chemical contains too much iron (III), to be really useful. Ferrous sulfate does not keep well. The pure compound, which is light blue/green, quickly becomes covered by a greenish/brown crust, due to oxidation by oxygen from the air. Solutions of ferrous sulfate are very prone to aerial oxidation. Ferrous sulfate, technical grade, is available from pottery and ceramics shops and from photography raw chemical suppliers. The pure stuff is available from chemical supply houses, but the purchase of the latter is not recommended due to its bad storage properties.
A much better alternative for purchasing ferrous sulfate is the purchase of ferrous ammonium sulfate. This salt, also known as Mohr's salt, keeps quite well, provided it is stored in a dry place. It is light blue/green. Its solutions also are prone to aerial oxidation, but not as strong as the solutions of ferrous sulfate. The slight acidity, introduced by the ammonium ions, provides a fairly good protection for the ferrous ions against aerial oxidation. Mohr's salt is available from photography raw chemical suppliers as special order item and from chemical supply houses.
A nice source for ferric ions is ferric ammonium sulfate or ferric sulfate. The solutions of these compounds are light brown/yellow and somewhat turbid, due to hydrolysis of the iron (III) ions. Both salts are available from photography raw chemical suppliers. Ferric ammonium sulfate dissolves in water easily. Although ferric sulfate is more soluble than ferric ammonium sulfate, the dry salt dissolves in water very slowly. Dissolving a spatula full of this in some water may take many hours. Ferric ammonium sulfate is a crystalline solid, consisting of large transparent crystals, which are pale violet/brown. Ferric sulfate is sold as a dry compact powder, which is off-white.
Another alternative for ferric ions is ferric chloride. This dissolves in water very well. It is sold in the form of yellow/brown beads, or as a concentrated solution in water. The solid is deliquescent. It attracts water from the air very easily. When ferric chloride is used, then the iron (III) ions are not available as free ions, but coordinated to chloride ions. The chloro-complex of iron (III) is yellow/brown. Formation of this complex can be observed nicely by adding a solution of table salt to a solution of ferric sulfate or ferric ammonium sulfate. Ferric chloride is available from electronics shops as a PCB-etchant.
Ferric oxide is available from pottery and ceramics shops as a dark red/brown crystalline powder. This crystalline solid is calcined ferric oxide. This calcined oxide is quite inert. When some of this powder is added to concentrated hydrochloric acid, then even after boiling for several minutes, only trace amounts have dissolved. The calcined compound is so inert, that it hardly is interesting for the home chemist. A nice, but dangerous, dry-chemical experiment, which can be conducted with the finely powdered ferric oxide is the so called 'thermite' reaction with finely powdered aluminium metal. Beware, temperatures reached in this reaction become ultra-high, 2000+ °C and molten iron can be sprayed around!!
Ferrous sulfide is sold in the form of black lumps. It is not really interesting for the purpose of studying iron chemistry, but it can be interesting for making hydrogen sulfide gas with dilute hydrochloric acid. Making hydrogen sulfide with ferrous sulfide is more convenient than making hydrogen sulfide with sodium sulfide. It allows one to generate hydrogen sulfide at a more constant rate. Ferrous sulfide is available from chemical supply houses. It also can be made fairly easily by igniting a stoichiometric mixture of sulphur and iron powder. This should be done outside on a concrete tile.
Potassium ferrocyanide and potassium ferricyanide are used extensively in photography and alternative processes. The ferricyanide can be purchased at almost every photography raw chemical supplier. The ferrocyanide also is available at many of these suppliers, certainly as a special order. These chemicals form brightly colored precipitates with many metal-salts. This property is exploited in photography in many toner solutions. On this site, a treatise on ferrocyanide toners is given in the photography section. Both of these chemicals are a valuable addition to the home lab. These compounds do not contain free iron ions and they also do not contain free cyanide ions. The cyanide ions are bound to the iron-core very tightly. These compounds are not very toxic. Even in dilute acids, the cyanide is not released easily. Experimenting with these chemicals is quite safe. Only in hot or concentrated acids, one must take into account the possible release of hydrogen cyanide.
The "ferric ammonium citrates" also are used extensively in photography and alternative processes. These compounds also are used in medicine for treating iron-deficiency. Ferric ammonium citrate is not a well defined compound of fixed stoichiometry. It can be regarded as a compound, containing ferric ion, ammonium ion, free ammonia, and citrate ion at variable ratio. The citrate ions at least partially are coordinated to the ferric ions. The green compound has a lower iron content and has a somewhat higher ammonium content than the brown compound. It is slightly more acidic in its solutions. "Ferric ammonium citrate" is not the first iron compound to start with. If one wants to study the chemistry of iron, then the sulfates and to a lesser extent the chloride are more suitable. The presence of the citrate ion has a strong influence on the observed chemistry.
The "ferric oxalates", like the citrates, also are used extensively in photography and alternative processes. "Ferric ammonium oxalate" is a well defined compound, which, however, does not contain free ferric ions. The correct name for this compound is ammonium trisoxalato-ferrate (III) and in aqueous solutions it contains Fe(C2O4)33- ions. "Ferric oxalate" is not a well-defined compound. The most pure compound can be regarded as ferric trisoxalato-ferrate (III), but the commercial preparations usually contain appreciable amounts of oxalic acid. As the citrates, these compounds are not the most suitable when one wants to study the properties of iron in its aqueous solutions. The chemistry of trisoxalato ferrate (III) differs considerably from the chemistry of free iron (III) and free oxalate.