Iron pentacarbonyl
| Names | |
|---|---|
| IUPAC name
Pentacarbonyliron(0)
| |
| Other names
Iron carbonyl
Pentacarbonyl iron | |
| Properties | |
| Fe(CO)5 | |
| Molar mass | 195.90 g/mol |
| Appearance | Straw-yellow or orange liquid |
| Odor | Musty, stinky rubber |
| Density | 1.453 g/cm3 (20 °C) |
| Melting point | −21 °C (−6 °F; 252 K) |
| Boiling point | 103 °C (217 °F; 376 K) |
| Insoluble | |
| Solubility | Miscible with acetic acid, benzaldehyde, bromobenzene, carbon disulfide, CCl4, ethyl acetate, pentanol, tetralin Soluble in acetaldehyde, acetone, benzene Slightly soluble in alcohol Insoluble in liq. ammonia |
| Vapor pressure | 40 mmHg (30.6 °C) |
| Thermochemistry | |
| Std molar
entropy (S |
338 J·mol-1·K-1 |
| Std enthalpy of
formation (ΔfH |
-764 kJ/mol |
| Hazards | |
| Safety data sheet | Sigma-Aldrich |
| Flash point | −15 °C (5 °F; 258 K) |
| Lethal dose or concentration (LD, LC): | |
| LD50 (Median dose)
|
25 mg/kg (rat, oral) |
| Related compounds | |
| Related compounds
|
Nickel tetracarbonyl |
| Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa). | |
| Infobox references | |
Iron pentacarbonyl, also known as iron carbonyl, is the compound with formula Fe(CO)5. Iron pentacarbonyl is a homoleptic metal carbonyl, where carbon monoxide is the only ligand complexed with a metal.
Contents
Properties
Chemical
Iron pentacarbonyl will slowly decompose to yield iron tetracarbonyl or iron nonacarbonyl, when exposed to light (photolysis). UV light will cause it to decompose even faster. This can be easily seen by exposing some (FeCO)5 to a UV lamp or a blue/purple laser.
Iron pentacarbonyl burns when ignited, but it does not produce a clear visible flame, instead copious amounts of iron oxide fumes are being produced. The flame can be seen better is liquid oxygen is added. Adding a peroxide, like tert-butyl hydroperoxide will cause a delayed reaction, resulting in ignition. However, some oxidizers, like liq. nitrous oxide are added, it still does not produce a visible flame. Meaning that the compound tends to decompose rather than burn.[1]
Iron carbonyl is a reducing agent, capable of reducing iron(III) salts to iron(II), as well as silver salts to their respective metals.
Treatment of Fe(CO)5 with iodine gives iron tetracarbonyl diiodide:
- Fe(CO)5 + I2 → Fe(CO)4I2 + CO
Iron pentacarbonyl is not readily protonated, but it is attacked by hydroxides. Treatment of Fe(CO)5 with aqueous base produces Potassium tetracarbonyliron hydride ([HFe(CO)4]−), via the metallacarboxylate intermediate.
Physical
Under standard conditions, iron pentacarbonyl is a free-flowing, yellow-orange-colored liquid, although when frozen it turns completely yellow.[2] It has a pungent odor, sometimes described as "stinky rubber". Older samples appear darker. It is insoluble in water, but it is more soluble in other organic solvents.
Availability
Iron pentacarbonyl is sold by chemical suppliers, though it cannot be acquired by private individuals due to its hazards.
Preparation
Iron pentacarbonyl can be synthesized by reacting carbon monoxide with iron(II) iodide, then the resulting product is reduced using copper:[3]
- FeI2 + 4CO → Fe(CO)4I2
- 5 Fe(CO)4I2 + 10 Cu → 10 CuI + 4 Fe(CO)5 + Fe
Iron pentacarbonyl can be produced by the reaction of fine iron particles with carbon monoxide, at high pressure and temperature.[4]
Projects
- Make adducts
- Preparation of carbonyl iron
- Make iron ball paint (radar absorbent materials)
- Flame speed inhibitor
Handling
Safety
Iron pentacarbonyl is toxic, which is of concern because of its significant volatility. If inhaled, iron pentacarbonyl may cause lung irritation, toxic pneumonitis, or pulmonary edema.
Decomposition of Fe(CO)5 will yield carbon monoxide, which is extremely toxic if inhaled.
Storage
Iron pentacarbonyl should be kept in airtight bottles, away from any light and heat. The container should be wrapped in cloth or some other absorbent material.
Disposal
Should be burned outside.
References
- ↑ http://www.youtube.com/watch?v=CRuUA0ytsjc&t=7m40s
- ↑ https://www.youtube.com/watch?v=CRuUA0ytsjc
- ↑ Brauer, Georg (1963). Handbook of Preparative Inorganic Chemistry. Vol. 2 (2nd ed.). New York: Academic Press. pp. 1743, 1751
- ↑ https://zenodo.org/record/1702462#.YxSjhEZBzIU
