Anhydrous Dinitrogen Trioxide Solutions for Br?nsted Acid Free Nitrous Acid Chemistry

Article abstract of DOI:10.1002/ejic.201701091

Dinitrogen trioxide, N₂O₃, is readily prepared and stabilized in high concentrations in dry organic solvents at normal working temperatures. These conditions allow for facile acid and water free nitrosation and nitration reactions fo

Full text of DOI:10.1002/ejic.201701091

DOI: 10.1002/ejic.201701091
Full Paper
Nitrous Acid Chemistry
Anhydrous Dinitrogen Trioxide Solutions for Brønsted Acid Free
Nitrous Acid Chemistry
Kristopher A. Rosadiuk[a] and D. Scott Bohle*[a]
Abstract: Dinitrogen trioxide, N2O3, is readily prepared and sta- three examples are given here: the preparation of benzene-
bilized in high concentrations in dry organic solvents at normal diazonium nitrite, [PhN2][NO2]; nitrosyl chloride, NOCl, and
working temperatures. These conditions allow for facile acid nitrosylsulfuric acid, (ONOSO3H).
and water free nitrosation and nitration reactions for which
Introduction
Vosper reports a few chemical experiments in which these solu-
tions are treated with primary and secondary amines, and some
metals. Yet the stabilizing effect of organic solutions on N2O3
remains largely underutilized, and although re-discovered or
noted obliquely from time to time,[5–7] it has not been further
developed.
In our laboratory, we have found that organic solvent stabi-
lized N2O3 can be easily prepared at molar concentrations with
reliable purity, often superior to prior reports and without the
use of cryogens. Furthermore, it may be stored and used at
room temperature, or higher, to perform a wide variety of dif-
ferent chemical operations, from nitration/nitrosation to oxid-
ation and adduct formation. In this paper we describe our
methods for the preparation and handling of these solutions,
as well as three original examples of convenient uses for the
solutions in the preparation of otherwise tricky derivatives:
benzenediazonium nitrite, pure nitrosylsulfuric acid, and
nitrosyl chloride.
Nitrosation and nitration are abiding challenges for green
chemistry. The strongly acidic and oxidizing environment of
conventional nitrous acid mediated reactions is problematic for
many substrates and can engender a host of by-products. At
the heart of this chemistry is dinitrogen trioxide (N2O3), an
oxide of nitrogen that reversibly condenses as an azure blue
liquid when the radicals nitric oxide (NO) and nitrogen dioxide
(NO2) are free to combine at temperatures below –40 °C.[1]
Nitric oxide and dinitrogen trioxide are also termed nitrogen
monoxide and nitrogen trioxide; here the former will be used
without radical dots. It dissociates readily upon warming and
as such has been little studied, at least compared to its constitu-
ent gas species or dinitrogen tetroxide. Most of what is known
of its chemistry is in connection with nitrous acid, which is its
hydrated form [Equation (1)].
(1)
The Keq of this reaction in water is estimated[2] to be
Results and Discussion
3 × 10–3
M
–1, which is sufficient to lend a pale blue coloration
to concentrated nitrous acid.
N2O3 in Organic Solvents
Nevertheless, it has been known for over a century that di-
nitrogen trioxide can be stabilized by dissolving it in an organic
matrix, provided that the solvent is either polar or aromatic
with no active hydrogens. The earliest report of such an effect,
to the best of our knowledge, is from Cundall[3] in 1895, who
dissolved it in various chlorinated solvents. Joan Mason spectro-
scopically characterized dissolved N2O3 at low temperatures,[4]
but the most recent and thorough study of the solutions
themselves has been the work of A. J. Vosper, who published a
series of papers from 1960 to 1974, which describe how
solutions of N2O3 can remain stable even at room temperature.
Shaw and Vosper report[8] that nearly pure N2O3 can be pre-
pared by condensing NO2 gas as solid N2O4 at –78 °C in the
upper regions of a long vessel submerged in a dry ice/ethanol
bath, then passing NO gas over this such that liquid N2O3 may
condense as it forms, pooling in the bottom of the vessel. The
solvent is added afterwards. Although this is effective it is a
cumbersome operation that limits the throughput of the prepa-
ration and requires extensive use of cryogens.
We have found that a simple and reliable alternative is to
use a system of dynamic absorption, in which an ice-chilled
solvent is treated with small alternating doses of oxygen and
NO while under gentle stirring (see Experimental for details).
NO gas reacts with oxygen to produce NO2, which combines
with the next round of NO at the surface of the solution to
form N2O3. The N2O3 dissolves rapidly, allowing more gas to be
introduced. In this way the chemist can add over a liter of gas
[a] Department of Chemistry, McGill University,
801 Sherbrooke St.W., Montreal H3A 0B8, Canada
E-mail: Scott.bohle@mcgill.ca
http://bohle-group.mcgill.ca
ORCID(s) from the author(s) for this article is/are available on the WWW
under https://doi.org/10.1002/ejic.201701091.
Eur. J. Inorg. Chem. 2017, 5461–5465
5461
© 2017 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim

Full Paper
to only a few milliliters of solvent without the application of phere that is not already suffused with NO gas will cause some
high pressure. For proper use, it is crucial to understand that dinitrogen trioxide degradation as the more volatile NO compo-
solutions of N2O3 are in a state of equilibrium with their atmos- nent is driven out of solution.
phere [Equation (2)].
The capacity of various solvents to maintain the N2O3/NO–
NO2 equilibrium varies with their polarity (greater polarity
allows greater N2O3 content) and with the electron donating
character of the aromatic ring in aromatic solvents. Other fac-
tors such as viscosity and boiling point can make a difference,
but the electronic character predominates, as it stabilizes the
N–N bond and prevents the N2O3 from dissociating. Table 2 lists
some common solvents and the concentration of N2O3 that
may be achieved in them through the dynamic absorption
method under the conditions described below (the balance of
the mol-% will be N2O4, which is also stabilized by these sol-
vents[10]).
(2)
In the gas phase, the equilibrium constant favors dissociation
to NO and NO2, at 25 °C, Keq = 1.916,[9] while the presence of
solvent shifts the equilibrium to the left. Nevertheless, the in-
tegrity of the solution is only as good as the atmosphere of NO
above it, which is why Shaw and Vosper specify a high excess
of NO. Losses of NO in the atmosphere will drive Equation (2)
forward, converting the blue-green N2O3 solution to yellow di-
nitrogen tetroxide (N2O4). Higher temperatures also drive the
reaction forward. In many solvents, a 1 atm partial pressure is
Based on our titration experiments, we find that our system
of absorbing NO directly into the solvent system allows for far
higher concentrations to be formed and maintained than what
Vosper initially reported, even at elevated temperatures. In his
research, Vosper attempted to prepare N2O3 by absorbing NO
gas directly into liquid N2O4, with violent agitation and without
another stabilizing solvent. With this method at –5 °C the maxi-
mum N2O3 concentration was 54 %.[8] Furthermore, it was
concluded that at 0 °C it was impractical to maintain an N2O3/
N2O4 solution beyond around a 60:40 ratio[12] In this case it was
prepared pure at –78 °C and then warmed with concentrations
being, based on the UV/Vis spectrum, which has significant
overlap between the two compounds. We believe that our
measurement method is more robust and reflects the true po-
tential concentrations.
We confirm prior results that almost any solvent that is either
(a) polar and/or (b) aromatic can serve to stabilize N2O3, with
the exception of protic solvents such as water or alcohol, which
react with it. Our laboratory has thus far worked with stable
solutions in acetic anhydride, acetone, acetonitrile, benzalde-
hyde, benzene, chlorobenzene, chloroform, dichloromethane,
dimethyl sulfoxide, ethyl acetate, mesitylene, 1,1,2,2-tetrachlo-
roethane, tetrahydrofuran, toluene, and xylene. It is possible to
form a solution in anisole, but reaction with the solvent is ap-
parent within an hour; benzaldehyde will also react (forming
benzoic acid), though over a span of days.
sufficient to maintain a concentration of about 1
in near purity.
M N2O3 at 0 °C
It is thus useful to think of N2O3 solutions the way one thinks
of carbonated beverages, which maintain their concentration of
carbonic acid so long as they are capped. When opened, this
carbonic acid will steadily convert back to carbon dioxide and
bubble out, gradually causing the beverage to go “flat”. Like
carbonated water, organic N2O3 solutions takes some time to
degas, and a beaker of blue solution can be kept in the open
air for several minutes before losing considerable amounts of
N2O3. Table 1 lists four typical solvents and the measured de-
gassing rate from a surface area of 1 cm2 (for a vessel of 2.7 mL
volume, kept inside a 4 L container in between tests to limit
water absorption and air circulation). Thus a solution in toluene
left without agitation at room temperature and protected from
air currents will have approximately 1/2 of its starting N2O3 con-
centration after 2.6 hours.
Table 1. Rate of dinitrogen trioxide loss from a still solution.[a]
Solvent
kloss [s–1
]
τ1/2 [h]
Acetonitrile
Dichloromethane
Toluene
1.9 × 10–4
3.8 × 10–4
7.3 × 10–5
5.2 × 10–5
1.5
0.8
3.8
5.3
Chlorobenzene
[a] 1 cm2 surface at 25 °C. Estimated errors are in the last digit.
It is thus preferable to work with dinitrogen trioxide solu-
The exact mechanism(s) for the stabilization in these solu-
tions in a “one-pot” manner, in which the solution is first tions remains speculative. Prior workers have proposed that the
formed and further reactants are then added to the same ves- stabilizing mechanism in aromatic solvents was due to a charge
sel. Transfer of the dinitrogen trioxide solutions into an atmos- transfer interaction.[13] Low temperature studies of N2O4 dis-
Table 2. N2O3 content of solutions prepared by dynamic absorption.
Solvent
mol-%
mol-%
ε (25 °C)
ε (0 °C)[a]
N2O3 (0 °C)
N2O3 (25 °C)
[L/cm mol]
[L/cm mol]
Acetonitrile
Dichloromethane
Chlorobenzene
Benzene
Toluene
Xylene
91.7(6)
76.8(5)
91.5(7)
88.7(4)
99.8(7)
98.5(8)
100.7(7)
73.6(5)
59.1(4)
55.7(4)
66.5(3)
69.0(3)
69.0(5)
71.1(5)
9.8(7)
8.5(7)
10.7
9.2
10.9(9)
14.0(8)
15.3(7)
17.4(8)
16.4(8)
15.2
16.3
18.7
20.3
20.8
Mesitylene
[a] ε values taken from Shaw and Vosper,[11] which were determined at lower temperatures. These values were not found to hold at 25 °C and were thus
recalculated, but are accurate up to 0 °C.
Eur. J. Inorg. Chem. 2017, 5461–5465
www.eurjic.org
5462
© 2017 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim

Full Paper
solved in benzene[14] have shown that orbital overlap between the preferred solvent system employed should be immiscible
the dinitrogen tetroxide and the aromatic rings produces a kind with water and the product appears after twenty minutes as a
of sandwich compound, so it is probable that a similar interac- brown-orange suspension. The product is readily isolated by
tion goes on in N2O3 (although the color change of N2O4 in passing a stream of nitrogen through the solution and condens-
aromatic solvents[10] can be more pronounced than what is ing the nitrosyl chloride as a bright orange solid with a liquid
seen in N2O3). Polar solvents stabilize N2O3 as well, but to a nitrogen cold trap. This use of a reagent in a biphasic reaction
lesser degree; this is likely because this environment favors the minimizes the amount of hydrolysis of the product and gives a
more polar N2O3 molecule (2.572 D[15]) over its less polar sub- 45 % yield for the crude and 25 % yield for the final purified
stituents (NO2 0.289 D;[16] NO 0.160 D[17]), making it the enthal- product.
pically favored form.
(5)
Three Examples of Anhydrous N2O3 Chemistry
As reported,[18] Primary and secondary amines[18] react with
Conclusion
N2O3 solutions similar to nitrous acid, an unsurprising fact given
Though considered a laboratory curiosity and infrequently stud-
that N2O3 is believed to be the reactive form of HONO2, behav-
ied in the past century, N2O3 is an important nitrogen oxide
that plays a role as an intermediate in many nitrogen reactions.
By stabilizing and concentrating this intermediate under normal
laboratory conditions, we can transform an esoteric substance
into a common laboratory reagent, allowing both an easier
study of its chemistry and access to useful processes. Merely by
eliminating water as the solvent in standard nitrous acid reac-
tions, we can more easily produce and isolate pure nitrosation
products. More exotic chemistry also becomes possible by reac-
tion with aprotic systems and by the movement into a new
temperature regime (from extreme cold, where most N2O3
chemistry is performed, to room temperature and beyond). It is
our hope that N2O3 will come to be regarded as a useful rea-
gent and that these solutions will become a part of the
chemist's expanded toolbox. Future publications will report on
the novel chemistry that is possible by treating tertiary amines
with anhydrous N2O3 solutions.
ing as an [NO+] donor. In general, any system in which protons
may be transferred to N2O3 can be considered analogous to
nitrous acid,[19] so it is in its aprotic reactions under anhydrous
conditions that the chemistry of N2O3 solutions becomes
unique. The transition from aqueous to non-aqueous solvents
opens up some interesting possibilities; reactions that normally
make use of HNO2 can be conducted in water-immiscible sol-
vents to improve separation, or even to allow for the isolation
of hydrolytically sensitive products that would otherwise be de-
stroyed in water. Three laboratory preparations for which an
N2O3 preparation offers advantages over nitrous acid are shown
in Equations (3), (4), and (5) and are detailed in the Experimen-
tal Section.
Diazonium cations are often intermediates in aromatic sub-
stitution chemistry and are seldom isolated. However under cer-
tain conditions with the inert anions such as tetrafluoroborate
stable salts are accessible after metathesis. Unfortunately the
nitrite/nitrate salts of benzenediazonium, for example, are noto-
riously unstable[20] and must usually be prepared with special
measures in place to remove excess nitrous acid. We find how-
ever these salts can rapidly be prepared in high yield and purity
using anhydrous N2O3 system, Equation (3), from an aceto-
nitrile/benzene solution. The off white salts can be filtered and
used as is or recrystallized from benzene.
Experimental Section
Preparation of N2O3 Solutions
A sealed 50 mL Schlenk flask equipped with a stir bar is evacuated
under high vacuum. It is then flushed with nitric oxide gas. Between
10 and 25 mL of the chosen solvent is then injected with a syringe.
Alternatively, solvent can be added first, then freeze-thaw-degassed
or sparged with argon, and then finally flushed with NO (if it is not
necessary to know the precise quantity of N2O3 being produced,
the flask with a dry solvent can simply be flushed with NO gas;
residual oxygen will result in a small amount of extra N2O3 being
produced, but this will be inconsequential for many experiments).
The flask with its solvent and NO atmosphere is then immersed in
an ice-water bath and cooled to 0 °C. A measured amount of oxy-
gen gas is then injected in increments of ca. 10 mL with a gas tight
syringe. Each O2 addition is followed by allowing the mixture to stir
gently for ca. 5 min while under a weak positive pressure of NO gas
(ca. 0.7 kPa above atmospheric pressure), after which the next dose
of O2 may be added. NO and O2 combine to form NO2, which con-
sumes another NO to make N2O3, which dissolves in solution. Note
that each mol of O2 forms two mol of N2O3.
(3)
In the case of nitrosylsulfuric acid, also recognized as cham-
ber crystals from the chamber process, a pure powder prepara-
tion is not normally possible except from a fuming HNO3/SO2
gas reaction. With dinitrogen trioxide it can be prepared di-
rectly from laboratory grade sulfuric acid in acetonitrile [Equa-
tion (4)]. The nitrosylsulfuric acid product is a fluffy white solid
which can be isolated and used at this stage or recrystallized
from acetonitrile/dichloromethane.
The resulting solution of N2O3, when kept at this temperature or
below and with an atmosphere of NO over top of it, will keep indefi-
nitely, though in practice it is best to use it immediately. Long term
(4)
Finally nitrosyl chloride can also be prepared by treating an-
hydrous N2O3 with concentrated HCl [Equation (5)]. In this case storage requires the use of glass stoppers, as rubber septa are de-
Eur. J. Inorg. Chem. 2017, 5461–5465
www.eurjic.org
5463 © 2017 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim

Full Paper
stroyed by the gas over a span of hours. The addition of liquids to The growth of an IR peak at 1386 cm–1 signals its formation. The
nitrogen trioxide solutions is best performed with a syringe through
a septum while the reaction vessel is allowed to vent to a gas bub-
bler or a gas collection vessel, as gaseous byproducts are common.
For precise measurements, added reactants should be cooled in the
appropriate medium before addition to a cooled trioxide solution;
failure to chill prior to addition can cause some nitric oxide/nitrogen
dioxide to boil out of solution (cold solutions must be added
promptly as pipette tips/syringe needles can condense atmospheric
water during transfer). Solids can be added quickly by means of a
solid addition funnel, and the vessel closed immediately afterwards.
products are otherwise very similar.
WARNING: Benzenediazonium nitrite/nitrate is highly explosive and
will detonate on warming to 80 °C, from sufficient friction, or from a
sharp blow. Synthesis in large quantities is not recommended, and
even small quantities should be filtered using gravity and a filter pa-
per; loosen the filtrate before drying. DO NOT allow to dry in a sintered
glass frit. It is soluble in H2O, EtOH, and MeOH.
Following a literature technique,[23] we attempted to ascertain the
purity of our product by means of chemical activity. The product
was treated with sodium thiosulfate in water, and the excess thio-
sulfate was titrated with an iodide/I2 solution, using a starch indica-
tor. The product reacts readily with thiosulfate in a 1:2 ratio to yield
an orange solution containing a range of aromatic products. As
determined by this method, the yield, in terms of functional
benzenediazonium from initial aniline reactant, is 76 %, which com-
pares well to the by-weight yield.
Determination of N2O3 Concentration: In order to arrive at the
values given in Table 2, the precise quantities of N2O3 in solution
were determined by titration with oxygen gas, which converts it to
N2O4, in the following manner: 10 mL of N2O3 solution was pre-
pared according to the dynamic absorption method described
above, at 0 °C using 10 mL of solvent and 30 mL of oxygen gas.
Once NO absorption was complete, the atmosphere inside the sol-
vent vessel was flushed out with a flow of argon (ca. 2–3 seconds
of strong flow, such that no new NO2 was seen to form when the
O2 is added). O2 titration was carried out by adding the gas incre-
mentally with a gas tight syringe (first by 10 mL, scaling down to
0.25 mL as the blue-green coloration vanished, which signaled an
approaching endpoint). At very low concentrations of N2O3, dissoci-
ation of NO into the headspace “hides” the remaining NO, so the
vessel was cooled in an ethanol/dry ice bath, which immediately
converts it back to visible N2O3. It was then warmed for 10 min and
titration was continued. When no blue-green was visible at –78 °C,
the final endpoint was determined by the UV/Vis absorption in the
660–720 nm region[21] (the peak position for residual N2O3 can vary
dramatically with the solvent used). It was compared to an N2O4
solution at the same temperature (N2O4 has a faint, broad absorp-
tion in this region and replaces the baseline). It has previously been
established by spectroscopic[12] and Evans-method NMR[19] tech-
niques that the N2O3 in solution is ca. 99 % in the covalent form,
rather than dissociated into [NO+][NO2–] or [NO][NO2].
Synthesis of Nitrosyl Chloride: One mL of concentrated aqueous
hydrochloric acid (36.5 %) is added to an equimolar dinitrogen tri-
oxide solution (20 mL at 0.5
M). In water immiscible solvents (i.e.
toluene, dichloromethane), 20 min of stirring allows for the forma-
tion of a cloudy yellow-brown solution; in acetonitrile, the solution
immediately becomes a clear red-orange. In any of these solutions
the main product is seen to be the same. Bubbling nitrogen gas
through the solution for 60 min and collecting the stream in a liquid
nitrogen cold trap gives nitrosyl chloride as a brilliant orange solid.
The functional yield in situ from an acetonitrile preparation was
determined to be 45 % by reacting the product with mercaptoetha-
nol and measuring the UV absorbance of the resulting nitrosothiol
(ε = 34.3 at 548 nm), assuming a 1:1 formation of nitrosothiol from
each nitrosyl chloride (thiols are extremely reactive with nitrosating
agents, while NOCl is unlikely to react with the ethanol compo-
nent[24]). Collected as a pure product in a cold trap, a 25 % yield
(based on the added HCl) can be recovered. Nitrosyl chloride evapo-
rates quickly at room temperature, but can be identified by the UV
spectrum of its vapor.[25] UV/Vis: εmax at about 220 nm, 341 nm,
smaller peaks at 435 nm, 474 nm.
Synthesis of Benzenediazonium Nitrite
A mixture of 5 mL of acetonitrile and 10 mL of benzene is cooled
Synthesis of Nitrosylsulfuric Acid: Sulfuric acid (0.5 mL) is intro-
duced to excess dinitrogen trioxide in acetonitrile (30 mL of 0.5
to 0 °C on an ice bath, and a 0.146
M solution of N2O3 is prepared
M
in it, as above, inside a 50 mL three necked flask. 0.1 mL of aniline
(distilled and recrystallized from ether at –20 °C, and stored over
molecular sieves) is diluted in 5 mL of benzene, and this solution is
added by pipette (at a rate of ca. 1 mL over 3 seconds) to the
mixture while the liquid is stirred smoothly and briskly. The blue of
the reaction mixture will disappear once the N2O3 is fully con-
sumed, and the solution becomes faintly yellow. Stirring is halted
and the white-to-yellow precipitate is allowed to form for 5 min. It
is then filtered by gravity on filter paper (see below), and washed
with cold benzene, then air dried (or vacuum dried on an open
watch glass) to yield glossy, off-white flakes of benzenediazonium
nitrite, pure by IR and NMR spectroscopy. The filtrate can be saved
and a second and sometimes third crop of product may form on
solution) to yield a white product which forms as dry granules. The
product will spontaneously decay to nitric oxide and sulfuric acid if
removed from the solvent. Instead, the product is washed three
times with cold dry acetonitrile, then dried under vacuum to yield
nitrosylsulfuric acid[26] (chamber crystals) in 78 % yield as a fluffy
white powder. Alternatively, washing with acetonitrile then stirring
with dry dichloromethane for 24 h will yield dense white chunks of
dry precipitate, in similar yield. The reaction can be performed in
other non-aromatic solvents such as dichloromethane, where it
forms as an oily paste that requires more drying. The final product
in any case decays in air with release of nitric oxide, and reacts
violently with water and alcohol, again releasing nitric oxide.
IR: ν = 3470 (br), 1656 (w), 1438, 1383, 1363, 1268, 1156 (str),
cooling. Yield by weight: 85 %. IR: ν = 3088, 3064, 3040, 2396 (w),
˜
˜
831 cm–1. Confirmed by comparison with product prepared by liter-
ature techniques[27] (absorption of NO gas into sulfuric acid), which
produces the 40 % by weight solutions (in sulfuric acid) normally
sold commercially.
2296 (w), 1756 (w), 1607, 1551, 1458 (s), 1358 (s), 1318 (s), 1310 (s),
1160, 1092, 830, 758 (s), 605 (s), 524 cm–1. NMR: The product reacts
slowly with all solvents that it dissolves in; the reaction with meth-
anol was sufficiently slow to perform NMR in CD3OD: 8.66 (d), 8.306
(t), 8.006 (t). No other impurities were detected, except for a trace
of the decay product, which grows in over ca. 24 h as the product
disappears.
Nitrosylsulfuric acid can be prepared in aromatic solvents, but this
produces intensely colored side products (red and orange) from
reaction with the solvent. In benzene the reaction is relatively clean
and this solvent may be used for the synthesis if necessary, though
the level of purity has not been ascertained. CAUTION: Use of
An excess of N2O3 during synthesis will instead tend to produce
benzenediazonium nitrate,[22] due to oxidation of the main product.
Eur. J. Inorg. Chem. 2017, 5461–5465
www.eurjic.org
5464
© 2017 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim

Full Paper
[10] C. C. Addison, Chem. Rev. 1980, 80, 21–39.
mesitylene as a solvent will produce a violent chain reaction, and
should be avoided.
[11] A. W. Shaw, A. J. Vosper, J. Chem. Soc., Dalton Trans. 1972, 961–964.
[12] A. J. Vosper, J. Chem. Soc. A 1966, 1759–1762.
[13] E. Niciecki, J. Vosper, J. Chem. Soc., Dalton Trans. 1978, 1721–1723.
[14] K. O. Stroemme, Acta Crystallogr., Sect. B 1968, 24, 1607–1614.
[15] S. Kishner, M. A. Whitehead, M. S. Gopinathan, J. Am. Chem. Soc. 1978,
100, 1365–1371.
Acknowledgments
Support from NSERC and the CRC for our research is gratefully
acknowledged.
[16] S. Heitz, R. Lampka, D. Weidauer, A. Hese, J. Chem. Phys. 1991, 94, 2532–
2535.
[17] Y. Liu, Y. Guo, J. Lin, G. Huang, C. Duan, F. Li, Mol. Phys. 2001, 99, 1457–
1461.
[18] D. L. Lovejoy, A. J. Vosper, J. Chem. Soc. A 1968, 2325–2328.
[19] K. Rosadiuk, McGill University, 2015.
Keywords: Nitrogen · Diazo compounds · Nitrosation ·
Nitration · Acidity
[20] H. Zollinger, Diazo Chemistry 1, VCH, Weinheim, 1994.
[21] J. Mason, J. Chem. Soc., Dalton Trans. 1975, 19–22.
[22] R. M. Scribner, J. Org. Chem. 1964, 29, 3429–3430.
[23] D. Ben-Efraim, The Chemistry of Diazonium and Diazo Groups; John
Wiley & Sons, New York, 1978.
[1] I. R. Beattie, A. J. Vosper, J. Chem. Soc. 1960, 4799–4802.
[2] D. Williams, Nitrosation Reactions and the Chemistry of Nitric Oxide, Elsev-
ier, Oxford, Amsterdam, 2004.
[24] D. T. Manning, H. A. Stansbury Jr., J. Am. Chem. Soc. 1959, 81, 4885–
4890.
[3] J. T. Cundall, J. Chem. Soc. Trans. 1895, 67, 794–811.
[4] J. Mason, J. Chem. Soc. 1959, 1288–1295.
[25] C. F. Goodeve, S. Katz, Proc. R. Soc. London Ser. A 1939, 172, 432–444.
[26] W. H. Schroeder, P. Urone, Environ. Sci. Technol. 1978, 12, 545–560.
[27] J. W. Mellor, A Comprehensive Treatise on Inorganic and Theoretical Chem-
istry, 1st ed., Longman, Green and Co. Ltd., London, 1928, vol. 8.
[5] T. Suzuki, T. Fukai, Y. Seki, M. Inukai, Chem. Pharm. Bull. 2009, 57, 89–91.
[6] L. Grossi, J. Strazzari, J. Org. Chem. 1999, 64, 8076–8079.
[7] P. Astolfi, M. Panagiotaki, C. Rizzoli, L. Greci, Org. Biomol. Chem. 2006, 4,
3282–3290.
[8] A. W. Shaw, A. J. Vosper, J. Chem. Soc. A 1970, 2193–2195.
[9] I. R. Beattie, S. W. Bell, J. Chem. Soc. 1957, 1681–1686.
Received: September 12, 2017
Eur. J. Inorg. Chem. 2017, 5461–5465
www.eurjic.org
5465
© 2017 Wiley-VCH Verlag GmbH & Co. KGaA, Weinheim