Nitrogen Triiodide

Simon Cotton
Uppingham School, Rutland, UK

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Nitrogen triiodide, NI3

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 (500 Kb).

2 NI3 (s) N2 (g) + 3 I2 (s)

before... ...after
Before......After

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 NI3.

Why is it so explosive?

The process 2 NI3 (s) N2 (g) + 3 I2 (s) is exothermic, so that N2 (g) + 3 I2 (s) 2 NI3 (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 NI3, an ammonia complex is obtained instead. This is either [NI3.NH3] or [NI3.(NH3)3], and the ammonia cannot be removed from this.

So can pure NI3 be made?

This wasn't achieved until 1990, when it was found that boron nitride reacted with iodine monofluoride in CFCl3 at -30°C.

BN + 3 IF BF3 + NI3

What is pure NI3 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. NCl3 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 NH4Cl. The main route used commercially today is the electrolysis of slightly acidic ammonium chloride; the NCl3 is removed as fast it is formed using an air current. This air/NCl3 mixture is much more stable than pure NCl3 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. NCl3 can be formed when chlorine reacts with nitrogen compounds in wastewater treatment plants. The particular danger associated with the formation of NCl3 under these conditions is that a combination of its sensitive nature and low solubility in water leads to explosive droplets of NCl3.

Stable NF3 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.

4NH3 + 3 F2 NF3 + 3 NH4F

NBr3 was originally synthesised in 1975 by the reaction of bis(trimethylsilyl)bromamine with ClBr at -78°C.

(Me3Si)2NBr + 2 BrCl NBr3 + 2 Me3SiCl

What are they like?

NF3 is pretty unreactive at room temperature; it is not affected by water and only reacts with most metals on heating. NCl3 is much more reactive; it is light-sensitive and, like all the other halides, apart from NF3, explosive. All these compounds are volatile, as expected for small covalent molecules.

Why is NF3 stable but the others are unstable?

For all these compounds, it is possible to work out DHf for the formation of NX3 in the gas phase, using bond energies.

N2 (g) + 3 X2 (g) 2 NX3 (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, DHf ) = - 123 kJ mol-1 (per mole of NF3); similarly, for ammonia, using H-H and N-H bond energies of 436 and 390 kJ mol-1, respectively, DHf = - 43 kJ mol-1 (per mole of NH3). In contrast, using I-I and N-I bond energies of 151 and 169 kJ mol-1, respectively, DHf per mole of NI3 = + 192 kJ mol-1.

Spacefill model of NI3One factor making NF3 more stable than the other NX3 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 NX3 molecules may be less stable than NF3 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 NI3, 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 109°; thus the value for ammonia, NH3, is 107.5°. Because fluorine is much more electronegative than hydrogen, the bond pairs of electrons are attracted away from nitrogen, so that in NF3 the bond angle is actually 102.3°. In contrast, the corresponding value for NCl3 is 107.1°, although on electronegativity grounds it would be expected to be intermediate between the values for NF3 and NH3. This may possibly be due to non-bonded Cl...Cl repulsions.

NF3 structure

The molecules themselves have a trigonal (triangular) pyramid shape. They are, of course, polar. NF3 has a small dipole moment (0.234D) in comparison with NH3 (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 NF3. NCl3 also has a small dipole moment (0.6D).

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

NCl3 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 NF3 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 NF3. 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 NX3
 NF3NCl3NBr3NI3
First synthesised 1928181119751990 (ammonia adduct 1813)
Description (at RT)Colourless gasPale yellow oilDeep red solidRed-black crystals
m.p. (°C) -208.5-40--
b.p. (°C) -129 71--
< X-N-X (°)102.3107.1--
N-X (pm)137175.9--
N-X bond energy (kJ mol-1)278188 169-
Dipole moment (D)0.2340.6--
DHf (kJ mol-1) -114232220 (est)159 (est)

Bibliography

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