# What is the most convenient way to prepare ferrous oxide (FeO) in the laboratory?

The Wikipedia page for ferrous oxide states that $$\ce{FeO}$$ can be prepared by the thermal decomposition of iron(II) oxalate, with the following reaction:

$$\ce{FeC2O4 → FeO + CO2 + CO}$$

And that the procedure is conducted under an inert atmosphere to avoid the formation of ferric oxide.

Sicius' Eisengruppe: Elemente der achten Nebengruppe: Eine Reise durch das Periodensystem states that $$\ce{FeC2O4}$$ should "slowly" be heated in a vacuum and then rapidly cooled down.

Both of these resources, however, omit critical information as to the preparation process. These questions remain:

• At which temperature does this reaction really start occurring, and what is the ideal temperature range to carry out this thermal decomposition?

• How "fast" should it be cooled down? And what are appropriate methods to do so?

I'd be very thankful for any additional information on the topic.

• If you were an undergrad student, this would qualify as "homework": literature research. Are you? ;-) – Karl Jul 28 at 20:11

In appears that the synthetic procedure for preparation of iron(II) oxide from iron(II) oxalate has been first described by Günther et al. [1] and subsequently summarized in Brauer's Handbook of Preparative Inorganic Chemistry [2, p. 1497]:

### Iron (II) Oxide

I.

$$\ce{\underset{143.8}{FeC2O4} = \underset{71.8}{FeO} + \underset{28.0}{CO} + \underset{44.0}{CO2}}$$

Thermal decomposition of $$\ce{FeC2O4}$$ yields pure $$\ce{FeO}$$ only under specific conditions. The decomposition is carried out in a quartz vessel (Fig. 332) whose lower section is kept at $$\pu{850 °C}$$ by means of an electric furnace. The joint is surrounded by a water-cooled lead coil or a rubber hose. The nascent gases should be removed as quickly as possible; for this reason, the reactor is connected to two parallel mercury pumps and a good forepump; the gas is carried into two liquid-nitrogen-cooled traps containing activated charcoal.

Fig. 332. Preparation of iron (II) oxide.

The starting $$\ce{FeC2O4}$$ $$(\pu{0.5 - 0.8 g.})$$ is placed in the small bulb above the quartz vessel, and the water of crystallization is completely vaporized by heating in vacuum for 12 hours at $$\pu{200 °C}.$$ The bulb is turned in the joint, and the $$\ce{FeC2O4}$$ drops into the heated lower section of the reactor where it is rapidly decomposed to $$\ce{FeO},$$ $$\ce{CO}$$ and $$\ce{CO2}$$ (the decomposition is complete in about 20 seconds). The product $$\ce{FeO}$$ is retained by a quartz wool plug, which must be loose enough to prevent a buildup of pressure during the decomposition.

The furnace is now removed and the hot quartz tube is chilled as rapidly as possible in cold water, since $$\ce{FeO}$$ is unstable in the range of $$\pu{300-560 °C}$$ and decomposes according to:

$$\ce{4 FeO = Fe3O4 + Fe}$$

(this decomposition proceeds most rapidly at about $$\pu{480 °C},$$ but ceases below $$\pu{300 °C}).$$ The above procedure yields a jet-black product, readily soluble in dilute acids; it is rapidly oxidized in air, but does not ignite.

Alternatively, iron(III) oxide can be reduced by iron to iron(II) oxide [2, p. 1498]:

II. The preparation from stoichiometric quantities of commercial $$\ce{Fe2O3}$$ and reduced iron can also be recommended. The mixture and a few drops of water are sealed into a preevacuated quartz tube, heated for about three days at $$\pu{900 °C},$$ and quenched in cold water.

$$\ce{Fe + Fe2O3 ->[\pu{900 °C}] 3 FeO}$$

Iron(II) oxide can also be produced from iron(III) oxide reduction by carbon monoxide at the temperature range between $$\pu{500 °C}$$ and $$\pu{600 °C}$$ (adapted from [3, p. 415]):

\begin{align} \ce{3 Fe2O3 + CO &->[\pu{400 °C}] 2 Fe3O4 + CO2} \\ \ce{Fe2O3 + CO &->[\pu{500-600 °C}] 2 FeO + CO2} \\ \ce{Fe2O3 + 3 CO &->[\pu{700 °C}] 2 Fe + 3 CO2} \end{align}

as well as by thermal decomposition of magnetite [3, p. 415]:

$$\ce{2 Fe3O4 ->[>\pu{1538 °C}] 6 FeO + O2}$$

### References

1. Günther, P. L.; Rehaag, H. Über Die Thermische Zersetzung von Oxalaten II. Mitteilung. Darstellung von Reinem Ferrooxyd. Z. Anorg. Allg. Chem. 1939, 243 (1), 60–68. https://doi.org/10/cqzgk8. (in German)
2. Handbook of Preparative Inorganic Chemistry, 2nd ed.; Brauer, G., Ed.; Academic Press: New York; London, 1965; Vol. 2. (Archive.org)
3. R. A. Lidin, V. A. Molochko, and L. L. Andreeva, Reactivity of Inorganic Substances, 3rd ed.; Khimia: Moscow, 2000. (in Russian)
• That's an awesome answer :) – Hans Jul 28 at 22:52

The paper linked below indicates that the proper temperature is north of 535°C. This decomposition should be carried out with great caution though as some of the products might enflame, or even initiate a thermite reaction if carried out in the presence of Aluminum (as in the experimental setup used for the paper hereafter).

• The CO would give one pause, unless this was being done with proper facilities and safety equipment. – Ed V Jul 28 at 21:06
• Agreed! Maybe with a way to vacuum extract the gases from the container once the decomposition has proceeded – Veritas Jul 28 at 21:11