Elemental Syntheses – Iron

Elemental Syntheses – Iron

December 12, 2019 5 By Bertrand Dibbert


In this installment of my Elemental Syntheses
series, I shall investigate the element iron. It’s a very common element, one of the few
recognised by the general public. However, the vast majority of iron seen is not pure;
and is usually better described as a steel – alloyed with a small amount of carbon, and
some other metals. However, mild or low-carbon steel is typically 0.05-0.25% carbon, with
a few other metals alloyed with the iron. Unless you’re after high purity samples for
your element collection, this is likely good enough. But I suspect you didn’t click on
this video just to learn that the core of the pennies in your pocket are fit for an
element collection, particularly with the thumbnail. I thought I’d do a slightly more
exciting preparation of iron. So, here is a traditional thermite: I lost the balls obtained in this thermite,
so I ran another one, but to no avail but these tiny spheres. This is probably because
of the scale I ran – at least 50 g and/or high purity ingredients are required for a
reliable thermite. All the same, that concludes the preparation section. Iron has two common oxidation states; iron(II)
trivially known as ferrous, and iron(III) trivially known as ferric. The trivial names
for oxidation states can get confusing; -ous is a lower oxidation state and -ic is a higher.
Not applicable for iron, but -ite is lower for an anion and -ate is higher. Hypo-…-ite
is lower than -ite but still anionic. Anyway, ferrous ions are typically green,
sometimes yellow, in colour and are relatively easily oxidised – air is often enough to oxidise
a sample to ferric. For this reason, the double salt ferrous ammonium sulphate is often used
as a source of ferrous ions when the ammonium would not interfere – this compound is more
stable to oxidation than ordinary ferrous sulphate.
Ferric ions are the oxidised form, and can be mildly oxidising, in contrast to ferrous
ions. The hydrated ions are red/orange/yellow in colour, and typically require mildly acidic
solutions to prevent hydrolysis and precipitation of hydrated iron(III) oxides. Interestingly,
if a non-coordinating counter-anion is used, for example ferric nitrate, in solid form
the compound is lilac, as shown here. The pictures are screenshots from a Periodic Videos
video, link in the description. Interestingly, once dissolved in water, the familiar orange
colour appears – this is due to the coordination and partial hydrolysis of the water. It is
this partial hydrolysis that means even solutions of ferric ions in distilled water are acidic.
There is one other oxidiation state of iron common enough to mention, and that is the
ferrate(VI) ion. (Side note: ferrite as an ion typically refers to iron(III) in the anion.
Ferrite can also refer to iron(II,III) oxide). Sodium ferrate can be prepared by heating
iron(III) oxide with a highly basic solution of bleach, sodium hypochlorite. The hypochlorite
oxidises the iron oxide to ferrate(VI), itself being reduced to chloride. An alternative
method, which is the route I took, is to pass chlorine into a very basic solution in which
iron oxide is suspended. I dissolved an arbitrarily large amount of sodium hydroxide in a test
tube full of water, and add a little iron oxide. Chlorine was then passed in, and the
solids allowed to settle. The result was a purple solution of sodium ferrate. This method
enables a higher concentration of sodium ferrate to be produced, because the chlorine is acting
as the oxidising agent, and is continuously being pumped in, rather than in the case of
sodium hypochlorite, where the concentration is fixed and rather low. Sodium or potassium
ferrate is an incredibly strong oxidising agent, stronger than the permanganate ion,
and the the byproduct of oxidation is iron oxide, rust. These properties have lent it
to being heavily researched as a clean and green oxidising agent. The only problem is
that it is only stable in very basic solutions, and even then decomposes slowly. All the same,
it is an unusual oxidation state of iron that is very easy to prepare at home – test tube
of bleach, a pinch of powdered rust, and a little sodium hydroxide. If no reaction occurs,
heat gently and wait for the purple colour to form. That’s all I’ve got for this video, thanks
for watching, particularly all the way to the end. Feel free to leave feedback in the
comments, or as a like, and I hope to see you in my next video!