Nitrogen Triiodide The well-known explosive crystals Simon Cotton

Uppingham School, Rutland, UK Molecule of the Month December 2001

Also available: JSmol versions.

Sounds like another boring covalent molecule

This one's explosive! When dry solid nitrogen triiodide is touched, even with a feather, it decomposes rather violently. Click on the images below to see an animation.

2 NI 3 (s) N 2 (g) + 3 I 2 (s)

Very impressive! The purple smoke is iodine vapour, I suppose?

Yes. In the animation, the shockwave from the first detonation sets off the second sample of NI 3 .

Why is it so explosive?

The process 2 NI 3 (s) N 2 (g) + 3 I 2 (s) is exothermic, so that N 2 (g) + 3 I 2 (s) 2 NI 3 (s) is endothermic. Endothermic compounds tend to be unstable.

Is that all?

Not at all, things are more complicated than they seem. Traditionally, nitrogen triiodide is made by reacting iodine with aqueous ammonia solution. That does not produce NI 3 , an ammonia complex is obtained instead. This is either [NI 3 .NH 3 ] or [NI 3 .(NH 3 ) 3 ], and the ammonia cannot be removed from this.

So can pure NI 3 be made?

This wasn't achieved until 1990, when it was found that boron nitride reacted with iodine monofluoride in CFCl 3 at -30°C.

BN + 3 IF BF 3 + NI 3

What is pure NI 3 like?

It's a dark red solid that can be sublimed in a vacuum at -20°C. It decomposes at 0°C, sometimes explosively.

Can you make other nitrogen trihalides?

Yes, certainly, thanks to some brave and intrepid chemists. NCl 3 was the first of the family to be made in 1811, by Pierre L.Dulong, who later became Professor of Chemistry at the Ecole Polytechnique in Paris; he lost 3 fingers and an eye in studying it. He made it by the reaction of chlorine with slightly acidic NH 4 Cl. The main route used commercially today is the electrolysis of slightly acidic ammonium chloride; the NCl 3 is removed as fast it is formed using an air current. This air/NCl 3 mixture is much more stable than pure NCl 3 and is commercially important. It is also formed in swimming pools when the chlorine gas used to disinfect the water reacts with nitrogen compounds found in urine, and can be a health risk to people like lifeguards who work continuously around the water. NCl 3 can be formed when chlorine reacts with nitrogen compounds in wastewater treatment plants. The particular danger associated with the formation of NCl 3 under these conditions is that a combination of its sensitive nature and low solubility in water leads to explosive droplets of NCl 3 .

Stable NF 3 was first made in 1928 by Otto Ruff, a German chemist (d.1939) who probably made more fluorides than anyone else, by electrolysis of a molten mixture of ammonium fluoride and hydrogen fluoride. Another route uses the reaction of ammonia with fluorine/nitrogen mixtures over a copper catalyst.

4NH 3 + 3 F 2 NF 3 + 3 NH 4 F

NBr 3 was originally synthesised in 1975 by the reaction of bis(trimethylsilyl)bromamine with ClBr at -78°C.

(Me 3 Si) 2 NBr + 2 BrCl NBr 3 + 2 Me 3 SiCl

What are they like?

NF 3 is pretty unreactive at room temperature; it is not affected by water and only reacts with most metals on heating. NCl 3 is much more reactive; it is light-sensitive and, like all the other halides, apart from NF 3 , explosive. All these compounds are volatile, as expected for small covalent molecules.

Why is NF 3 stable but the others are unstable?

For all these compounds, it is possible to work out ΔH f for the formation of NX 3 in the gas phase, using bond energies.

N 2 (g) + 3 X 2 (g) 2 NX 3 (g)

The process is not especially favourable owing to the difficulty in breaking the very strong N-N triple bond (E(N-N) = 945 kJ mol-1). Using values for the F-F and N-F bond energies of 159 and 278 kJ mol-1, respectively, ΔH f ) = - 123 kJ mol-1 (per mole of NF 3 ); similarly, for ammonia, using H-H and N-H bond energies of 436 and 390 kJ mol-1, respectively, ΔH f = - 43 kJ mol-1 (per mole of NH 3 ). In contrast, using I-I and N-I bond energies of 151 and 169 kJ mol-1, respectively, ΔH f per mole of NI 3 = + 192 kJ mol-1.

One factor making NF 3 more stable than the other NX 3 is the very low F-F bond energy (159 kJ mol-1). This has been ascribed to repulsions between lone pairs on the two rather proximate fluorine atoms. Additionally, the N-F bond is also particularly strong, as would be expected for a linkage between two elements in the first short period. The other NX 3 molecules may be less stable than NF 3 owing to congestion round the small central nitrogen atom leading to non-bonded repulsive interactions between the halogens. This is particularly bad for large iodine atoms, as can be seen in the space-fill image of NI 3 , right.

What is their structure?

In all these compounds, the nitrogen atom has a complete "octet", with four outer-shell electron pairs; one of these is a non-bonding ("lone") pair. These arrange themselves as far apart as possible around the nitrogen atom in a roughly tetrahedral disposition, to minimise repulsions between the negative charge clouds.

However, because repulsions involving lone pairs are stronger than those involving just bond pairs, the X-N-X angles are a little under the regular tetrahedral angle of 109½°; thus the value for ammonia, NH 3 , is 107.5°. Because fluorine is much more electronegative than hydrogen, the bond pairs of electrons are attracted away from nitrogen, so that in NF 3 the bond angle is actually 102.3°. In contrast, the corresponding value for NCl 3 is 107.1°, although on electronegativity grounds it would be expected to be intermediate between the values for NF 3 and NH 3 . This may possibly be due to non-bonded Cl...Cl repulsions.

The molecules themselves have a trigonal (triangular) pyramid shape. They are, of course, polar. NF 3 has a small dipole moment (0.234D) in comparison with NH 3 (1.42D); an explanation for this is that the moment due to the nitrogen atom and its lone pair is in opposition to the moment associated with the three polar N-F bonds in NF 3 . NCl 3 also has a small dipole moment (0.6D).

They sound very exotic- do they actually have any uses?

NCl 3 is used as a dilute mixture in air to bleach and sterilise flour and as a fungicide for citrus fruits and melons. The semiconductor industry uses NF 3 as an etchant of thin films, also for cleaning up chemical vapour deposition chambers, both uses depending on the use of a plasma to produce fluorine from NF 3 . It is also used as an oxidizer of high energy fuels, for the preparation of tetrafluorohydrazine (another fuel), and for the fluorination of fluorocarbon olefins, whilst it has been studies as a high-energy oxidiser for HF-DF chemical lasers.

A comparison of NX 3 NF 3 NCl 3 NBr 3 NI 3 First synthesised 1928 1811 1975 1990 (ammonia adduct 1813) Description (at RT) Colourless gas Pale yellow oil Deep red solid Red-black crystals m.p. (°C) -208.5 -40 - - b.p. (°C) -129 71 - - < X-N-X (°) 102.3 107.1 - - N-X (pm) 137 175.9 - - N-X bond energy (kJ mol-1) 278 188 169 - Dipole moment (D) 0.234 0.6 - - ΔH f (kJ mol-1) -114 232 220 (est) 159 (est)

Bibliography

N.N. Greenwood and A. Earnshaw, Chemistry of the Elements , Butterworth Heinemann, 2nd edition, 1997, 438-441

, Butterworth Heinemann, 2nd edition, 1997, 438-441 F.A. Cotton, C. Murillo, G. Wilkinson, M. Bochman and R. Grimes, Advanced Inorganic Chemistry , John Wiley, 6th edition 1999, 335-338

, John Wiley, 6th edition 1999, 335-338 J.E. Macintyre (ed), Dictionary of Inorganic Compounds , Chapman and Hall, London, 1992, entries IC-001954; IC-016907; IC-018291; IC-020275

, Chapman and Hall, London, 1992, entries IC-001954; IC-016907; IC-018291; IC-020275 H.H. Sisler, in Encyclopedia of Inorganic Chemistry , R.B. King ed, Wiley, 1994, Vol.5pp 2545-2551.

, R.B. King ed, Wiley, 1994, Vol.5pp 2545-2551. J. Jander, Adv. Inorg.Chem ., 1976, 19 , 1 (review of NX 3 , not X=F)

., 1976, , 1 (review of NX , not X=F) H.J. Emeléus, J.M. Shreeve and R.D. Verma, Adv. Inorg. Chem ., 1989, 33 , 139 (review of NF 3 )

., 1989, , 139 (review of NF ) Synthesis of NI 3 : I. Tornieporth-Oetting and T. Klapötke, Angew. Chem. Int. Ed. Engl. , 1990, 29 , 677

: I. Tornieporth-Oetting and T. Klapötke, , 1990, , 677 Synthesis of NBr 3 : J. Jander, J. Knackmuss and K-U. Thiedemann, Z. Naturforsch. Teil B , 1975, 30 , 464

: J. Jander, J. Knackmuss and K-U. Thiedemann, , 1975, , 464 Enthalpy of formation of NI 3 : R.H. Davies, A. Finch and P.N. Gates, J .Chem. Soc .Chem. Comm ., 1989, 1461.

: R.H. Davies, A. Finch and P.N. Gates, J ., 1989, 1461. NCl 3 structure : H. Hartl, J. Schöner, J. Jander and H. Schulz, Z. Anorg. Allgem. Chem ., 1975, 413 , 61

: H. Hartl, J. Schöner, J. Jander and H. Schulz, ., 1975, , 61 NI 3 adducts structure : J. Jander, L. Bayersdorfer and K. Höhne, Z. Anorg. Allgem. Chem ., 1968, 357 , 225.

: J. Jander, L. Bayersdorfer and K. Höhne, ., 1968, , 225. NCl 3 : http://inst.augie.edu/~djpaulso/indexA.html (rev)

: http://inst.augie.edu/~djpaulso/indexA.html (rev) Shape : R.J. Gillespie and I Hargittai, The VSEPR Model of Molecular Geometr y, Allyn and Bacon, Boston, 1991, p. 81.

: R.J. Gillespie and I Hargittai, y, Allyn and Bacon, Boston, 1991, p. 81. Stability : W.E. Dasent, Non-existent Compounds , Marcel Dekker, New York, 1965, pp 30-31

: W.E. Dasent, , Marcel Dekker, New York, 1965, pp 30-31 Uses and properties : http://www.semi.org/web/wsemi.nsf/364c709e8ffa8f9a882565de0080afe5/a666d06d06cb01ce8825696000815ef4!OpenDocument (NF 3 )

: http://www.semi.org/web/wsemi.nsf/364c709e8ffa8f9a882565de0080afe5/a666d06d06cb01ce8825696000815ef4!OpenDocument (NF ) http://www.unfccc.int/program/wam/wamsub024.html (NF 3 )

) http://www.c-f-c.com/gaslink/pure/nitrogen-trifluoride.htm (NF 3 )

) C. Shang, W-L. Gong and E.R. Blatchley, Environ.Sci.Technol ., 2000, 34 , 1721 (NCl 3 )

., 2000, , 1721 (NCl ) http://www.chem.purdue.edu/margerum/breakcl2.html (NCl 3 )

) http://www.amarillonet.com/stories/010497/010497.html (NCl 3 )

Back to Molecule of the Month page. [DOI:10.6084/m9.figshare.5245684]