A photochemical reaction of alkylammonium tetrachlorocuprates
by Kathleen Diane Mannila
A thesis submitted to the Graduate Faculty in partial fulfillment of the requirements for the degree of
MASTER OF SCIENCE in Chemistry
Montana State University
© Copyright by Kathleen Diane Mannila (1969)
Abstract:
A photochemical reaction has been observed in solutions of the alkylammonium tetrachlorocuprates,
and three mechanisms have been postulated. According to experimental results, none of the three is
precisely correct. It is believed, however, that copper (II) is reduced to copper(I) and that chloride is the
species oxidized. These observations point to the necessity of excluding stray light from solutions of
the tetrachlorocuprates prepared for quantitative spectral measurements, since a rapid decrease in
absorbtivity occurs even in ordinary laboratory illumination. 
In presenting this thesis in partial fulfillment of the require­
ments for .an .advanced degree at Montana State University, I agree that 
the Library shall make it freely available for inspection. I further 
agree that permission for extensive copying of this thesis for scholarly 
purposes may be granted by my,major professor, or, in his absence, by 
the Director of Libraries. It is understood that any copying or publica­
tion of this thesis for financial gain shall not be allowed without my 
written permission.
Signature: A
Date
A PHOTOCHEMICAL REACTION OF 
ALKYLAMMONIUM TETRACHLOROCUPRATES
by
KATHLEEN DIANE MANNILA
A thesis submitted to the Graduate Faculty in partial 
fulfillment of the requirements for the degree
of
MASTER OF SCIENCE 
in
Chemistry
Approved:
lead. Major Depattmetrf
Chairma Examining Committee
GraduateDean
MONTANA STATE UNIVERSITY 
Bozeman, Montana
December, 1969
-iii-
ACKNOWLEDGEMENT
The author gratefully acknowledges the advice and encouragement 
given by Dr. Kenneth Emerson, Dr. Paul W. Jennings, and Dr. Arnold C. 
Craig during the research on this project and during the writing of
the thesis.
TABLE OF CONTENTS
Page
LIST OF TABLES............................................................ v
LIST OF FIGURES... ....... ........................ ................. . vi
ABSTRACT............ ......... ...... ..................................... viii
INTRODUCTION................... ......... ....................... . I
DISCUSSION. . ................ .............................................  4
CONCLUSION... ..... ......... ......... ......... ...... ................ . 31
EXPERIMENTAL. ...... .............................. .................... .. 35
Preparation of Compounds................... ...... ............. 35
Spectral Measurements......... ................................ . 35
Decolonization of Tetrachlorocuprate Solutions With Light.... . 35
Determination of Light Region Causing Decolonization......... ... 36
Determination of Change in Free Chloride and pH. .............. . 36
Time Studies.................................................. . 37
LITERATURE CITED. ........ .......................................... . 42
-iv-
-V-
Page
I. Determination of Free Chloride........................      39
II= Determination of Change in Hydrogen Ion Concentration........... 39
III. Absorbance at 1200 m/*...................     40
Inp3IV. d Spacings for Precipitate Obtained from
Irradiation of Acetamidinium Tetrachlorocuprate................ . 41
LIST OF TABLES
-vi-
Page
1. Squashed Tetrahedral Structure of CuCl^ Ions in CSgCuCl^....... I
2. Square Planar Arrangement of CuCl^ Ions in (NH^)2re r 1^ ......... I
3. Visible and Near Infrared Spectra of Red Filter C o m i n g  2-59
and Blue-green Filter Corning 4-72........... ...................... 7
4. Ultraviolet Spectrum of Methylammonium Tetrachlorocuprate........ 8
5. Visible Spectrum of Methylammonium Tetrachlorcuprate.............  8
6. Ultraviolet Spectrum of Methylammonium Tetrachlorocuprate
After Decolorization be Light. . ................................... 9
7. Visible Spectrum of Methylammonium Tetrachlorocuprate
After Decolorization by Light...................... .............. .. 9
8. Ultraviolet Spectrum of Ethylammonium Tetrachlorocuprate......... 10
9. Visible Spectrum of Ethylammonium Tetrachlorocuprate.......... . 10
10. Ultraviolet Spectrum of Ethylammonium Tetrachlorocuprate
After Decolorization by Light............... ................. . 11
11. Visible Spectrum of Ethylammonium Tetrachlorocuprate
After Decolorization by Light................................. . 11
12. Ultraviolet Spectrum of Tetraethylammonium
Tetrachlorcuprate.................... ....... ........ ............ 12
13. Visible Spectrum of Tetraethylammonium Tetrachlorocuprate........ 12
14. Ultraviolet Spectrum of Tetraethylammonium
Tetrachlorocuprate After Decolorization by Light............. . 13
15. Visible Spectrum of Tetraethylammonium Tetrachlorocuprate
After Decolorization by Light...................................... 13
16. Near Infrared Spectrum of Methylammonium
Tetrachlorocuprate Before and After Decolorizatipn by Light.....  14
17. Near Infrared Spectrum of Ethylammonium 
Tetrachlorocuprate Before and After Decolorization by Light
LIST OF FIGURES
15
-vii-
18. Near Infrared Spectrum of Tetraethylammonium 
Tetrachlorocuprate Before and After Decolorization by Light...... 16
19. Ultraviolet Spectrum of Cuprous Chloride...................... .. „ . 18
20. Visible Spectrum of Cuprous Chloride............................. . 18
21. Near Infrared Spectrum of Cuprous Choride........................ . 19
22. Infrared Spectrum of ClCH CN + CuCl + (C R ), NCl
After Light.............. 7 ____....____ '. ......____ ______ _ 23
23. Infrared Spectrum of Chloroacetonitrile..... ................. . 24
24. Infrared Spectrum of Tetraethylammonium Chloride.... ............. 24
25. Infrared Spectrum of Tetraethylammonium
Tetrachlorocuprate Before Light................. .............. . 25
26. Infrared Spectrum of Tetraethylammonium
Tetrachlorocuprate After Light.......................... ....... . 25
27. Infrared Spectrum of Chlorine in Acetonitrile After Light..... . 28
28. Infrared Spectrum of Clg + CuCl + (CgH^^NCl After Light........ 28
29. Infrared Spectrum of Distillate from Tetrachlorocuprate
Irradiation.......... .............................................. .. 29
30. Visible Spectrum of Chlorine in Acetonitrile................ . 30
31. Decrease in Absorbance with Time. ............................ . 32
32. Optical Diagram of a Cary 14 Spectrophotometer. . „ . ............ 34
33. Standardization Curve for Determination of
Chloride Ion Concentration.......................................... 38
-viii-
ABSTRACT
A photochemical reaction has been observed in solutions of the 
alky!ammonium tetrachlorocuprates, and three mechanisms have been pos­
tulated. According to experimental results, none of the three is 
precisely correct. It is believed, however, that copper (II) is reduced 
to copper(I) and that chloride is the species oxidized. These observa­
tions point to the necessity of excluding stray light from solutions 
of the tetrachlorocuprates prepared for quantitative spectral measure­
ments, since a rapid decrease in absorbtivity occurs even in ordinary 
laboratory illumination.
INTRODUCTION
The existence of the tetrachlorocuprate (II) anion, CuCl^" ”, has been
reported for some time. A classification of the then known complex cupric
chlorides (33 in all) by Remy and Laves in 1933 showed 22 tetrachloro- 
8
cuprates. More recently, the study of the structure and spectra of cop­
per (II) halide complexes has received considerable attention.
Two different geometries have been observed for the CuCl^""ion. It
is known to exist as a distorted tetrahedron (Figure I) in CSgCuCl^,^ in
4
molten chlorides, and in acetonitrile and nitromethane solutions of
^^3^3^ 2^U^ 2 ^ U^^4*  ^ Spectral and X-ray meas­
urements indicate a square planar ion in (NH^)^CuCI^, in (CH^NH^)2CuCl^,
9 io
and in (C2H,.NH^^CuCI^ ’ (Figure 2). In general, MgCuCl^ compounds,
where M represents a univalent cation, contain nonplanar CuCl^ ions
3
provided the cations are large. Thus, monosubstituted ammonium salts 
contain square planar tetrachlorocuprate anions and tetrasubstituted 
ammonium salts contain tetrahedral anions.
Figure I. Squashed Tetrahedral 
Structure of CuCl^ ions in 
Cs2CuCI^ (a> p)
Figure 2. Square Planar
Arrangement of CuClf
4 9
ions in (NH^)2CuCI^
-2-
The electronic spectrum of the tetrachlorocuprate anion consists of 
a broad asymmetric band of medium intensity in the near infrared region, 
which can be identified as the d-d band of the central ion, and of a 
system of at least three intense bands in the visible and ultraviolet 
regions which can be assigned to a charge transfer mechanism operating 
from the halide ions to the copper(II) ion5 . The positions of the bands 
differ according to the geometry of the complex ion. Compounds contain­
ing tetrahedral CuCl^ are orange; hence their bands in the visible 
region occur at lower energy than do those of the light yellow compounds
containing square planar CuCl^ . Likewise, the d-d band appears at
-I
lower energy (6000-9000 cm ) in the tetrahedral species than in the
-I 9
square planar species (10,000-13,000 cm ).
During investigations of the spectral properties of complex ions
similar to the tetrachlorocuprate anion, several oxidation-reduction
reactions have been observed. Acrylonitrile solutions of copper(II)
chloro complexes ([cl 3: I) undergo a change in which Cu(II)
is reduced to Cu(I) and Cl is oxidized to elementary chlorine which
then reacts with the acrylonitrile. A similar reaction takes place
in solutions of copper (II) bromide. The reactions were carried out in
2
the absence of daylight and followed spectrophotometrically.
In a paper presented to the Seventh International Conference on 
Coordination Chemistry in 1962, Schneider stated that his spectral and 
magnetic studies of acetonitrile solutions of copper (II) bromo complexes 
were complicated by the formation of bromine and some copper(I) species.
-3“
No precipitation of CuBr occurred.7 Later, Furlani noted that although 
the frequency of the first charge transfer band of CuBr4 "" (16,500 cm"1) 
was sufficiently low to allow photochemical decomposition to take place 
easily, such decomposition did not occur so readily as to require special 
precautions during spectrophotometrie measurements.^
In related cases, however, such precautions have been shown to be 
necessary. Substitution reactions of platinum(IV) complexes are, for 
example, known to be light sensitive. The action of light on Ptdg"" 
leads to exchange with labeled halide ion in solution.
Ptci6"- + *cr  ^  P t V 6"  + or
The assumption is that the primary photochemical act is the formation 
of platinum(III) and chlorine atom.
PtClg"" + hkf -----> PtCl5"" + Cl*
A chain reaction could, then, be established.
PtCl5 + Cl ---— ►
*
Pt Cl5 + Cl"
Pt1tCl5"" + PtCl6""
*
Pt Cl6 + PtCl ""
The bromide complex, PtBr6 , behaves in a similar manner.1
Since photochemical reactions occur in the case of other chloro 
complexes, it is likely that they would also be present in solutions of 
the alkylammonium tetrachlorocuprates. When indications of this were 
observed, it was felt that an investigation of the characteristics and
extent of the reaction would be of considerable value.
-4-
DISCUSSION
Of the alkylammonium tetrachlorocuprates prepared in this laboratory, 
(CH^NH^)2CuCl^and (C2H^NHg)2CuCl^ were known to be square planar, while 
^ C2H5^4N^2CuCl4 Was k:nown to be tetrahedral. Another compound,
[CHgC(NHg)g]g CuCI^ was thought to exist as a mixture of the two geom­
etries . A decision was made to attempt to determine whether or not two 
structures were present and, also, to investigate the possibility of an
equilibrium between the two structures occurring in solution.
Since chlorocuprate structures had been previously determined from
-3
the positions of the spectral bands, spectra of 10 molar solutions of 
each of the available tetrachlorocuprates were taken in the near infra­
red, visible, and ultraviolet regions. At this time, an unexpected, 
inverted peak was observed at 1170 iq#. (See Figure 31) Because of its 
unusual nature, this portion of the spectrum was rerun after completion 
of the scan. The peak had not disappeared, but the absorbance at that 
wavelength had dropped by 0.1 a.u. (absorbance units). Later, the 
curious peak was shown to be an instrument artifact, leaving the question
of the decrease in absorptivity.
.
It soon became apparent that some type of photochemical reaction 
was taking place, and that the necessary energy was being supplied by the 
tungsten source. Solutions left in the sample compartment of the spec­
trophotometer dropped rapidly in absorbance; as much as 0.15 a.u. in 
ten minutes. Although the reaction continued to proceed when the cell 
was removed from the sample compartment and exposed to the light in
-5-
the laboratory, the rate, as measured by the decrease in the absorbance, 
was considerably reduced. Over a period of two days in room light, a 
drop of approximately 0.06 a.u. was observed.
Previous to these observations, spectral studies of copper halide 
complexes had been carried out using freshly prepared solutions. Stock 
solutions were not kept on the shelf indefinitely, as it was believed 
that some decomposition occured, possibly solvent exchange. However, a 
solution kept in a blackened flask and exposed to light only during 
measurement decreased 0.08 in absorbance over a period of twelve days. 
Hence, by far the most rapid decomposition occurred when the light inten­
sity was the greatest.
In a series of qualitative experiments, tetrachlorocuprate solutions 
were exposed to light from fluorescent tubes, tungsten bulbs, and direct 
sunlight. As would be expected, different reactions occured. The 
solution placed in fluorescent light acquired a greenish cast, and a 
finely divided precipitate formed; while both the solution placed in 
front of the tungsten lamp and that placed in the sunlight showed only a 
loss of color.
These results, as well as the behavior of the solutions in the 
spectrophotometer, seemed to indicate that the region of light causing 
reaction was the near infrared, possibly 800 to 1300 m/t, the position of 
the tetrachlorocuprate d-d band.
/ , .
To verify this, colored glass filters were placed between the sample
solution and the light source, a 500 watt tungsten projector lamp. Two
“6 “
filters were chosen: a red one which transmitted only light above 620nyt, 
and a blue-green one which allowed the passage of a band of light from 340
' I
to 600 nj4. Both filters excluded the ultraviolet portion of the spectrum 
(Figure 3). The solution placed in front of the red filter showed no 
visible color change or loss of color over a period of two and one half 
hours„ However, when the red filter was replaced by the blue-green, 
fading proceeded as before. This belied the initial postulate and pointed, 
instead, to the charge transfer band in the visible region.
It was next found that complete decolorization of the yellow tetra- 
chlorocuprate solutions could be effected by extended exposure to 500 
or 750 watt projector bulbs. The time required depended upon the con­
centration of the solution. The reaction could be speeded up through 
the use of mirrors or foil to reflect the light back through the sample.
It also seemed advantageous to cool the solution, as an increase in 
temperature appeared to slow down or even reverse the reaction.
For each of the tetrachlorocuprate solutions, (methy!ammonium-, 
ethy!ammonium-, and tetraethy!ammonium-) spectra were run before and 
after decolorization (Figures 4 through 18). In each case, the "after" 
spectrum showed an almost complete disappearance of the bands character­
istic of the tetrachlorocuprate ion. Of the solutions being investi­
gated, the acetamidinium compound was the only one showing any sign 
of precipitation in connection with the photochemical reaction. An 
attempt was made to analyze the precipitate by'X-ray diffraction, but 
the d values obtained did not compare with any possible species listed.
T
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4-72
WAVELENGTH
2-59
4-72 \
WAVELENGTH
Figure 3. Visible and Near Infrared Spectra of Red Filter Corning 2-59
and Blue-green Filter Corning 4-72
A
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Wavelength in mjj.
4. Ultraviolet Spectrum of Methylammonium Tetrachlorocuprate
- 0.9
- 0.8
-  0.6
- 0.4
-0.3
- 0.2
- 0.1
Wavelength in ir^u
Figure 5. Visible Spectrum of Methylammonium Tetrachlorocuprate
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Wavelength in np
6. Ultraviolet Spectrum of Methylammonium Tetrachlorcuprate 
After Decolorization by Light (0.1 cm. cell)
-0.9
-  0.8
-  0.6
- 0.3
-  0.2
- 0.1
Wavelength in np
Figure 7. Visible Spectrum of Methylammonium Tetrachlorocuprate after
Decolorization by Light (1.0 cm. cell)
-10-
- 0.90.9 -
-  0.80.8 -
0.7 - - 0.7
-  0.60.6 -
-0.50.5 -
_ 0.40.4 -
0.3 -
-  0.20.2 -
-  0.10.1 -
Wavelength in np
Figure 8. Ultraviolet Spectrum of Ethylammonium Tetrachlorocuprate
- 0.90.9 -
-  0.8
-0.70.7 -
- 0.6
~0.50.5 -
- 0.4
- 0.30.3 *
- 0.20.2 -
- 0.10.1 "
Wavelength in rtp
Figure 9. Visible Spectrum of Ethylammonium Tetrachlorocuprate
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- 0.9
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Wavelength in m»i
10. Ultraviolet Spectrum of Ethylammonium Tetrachlorocuprate 
after Decolorization by Light (0,1 cm. cell)
-0.9
-  0.8
- 0.6
- 0.4
- 0.3
- 0.1
Wavelength in nyi
Figure 11. Visible Spectrum of Ethylammonium Tetrachlorocuprate
after Decolorization by Light (1.0 cm. cell)
-12-
- 0.9
-  0.80.8 -
0.7 - -0.7
- 0.60.6 ~
0.5 -
_ 0.4
0.3 -
- 0.20.2 -
-  0.10.1 -
Wavelength in mfi
Figure 12. Ultraviolet Spectrum of Tetraethylammonium Tetrachlorocuprate
-0.90.9 -
-  0.80.8 -
-  0.60.6 -
0.5 '
- 0.4
- 0.30.3 *
- 0.20.2 -
- 0.10.1 -
Wavelength in np,
Figure 13. Visible Spectrum of Tetraethylammonium Tetrachlorocuprate
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- 0.9
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— 0.5
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Wavelength in np
14. Ultraviolet Spectrum of Tetraethylammonium Tetrachlorocuprate 
after Decolorization by Light (0,1 cm. cell)
0.9 - - 0.9
-  0.8
0.7 -
-  0.60.6 -
- 0.50.5 '
- 0.4
-0.3
0.2 -
- 0.10.1 -
Wavelength in m/u
Figure 15. Visible Spectrum of Tetraethylammonium Tetrachlorocuprate
after Decolorization by Light (10 cm. cell)
A
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-14-
0.5 *
0.4-
0 . 2 -
1000
Figure 16. Near Infrared Spectrum of Methylammonium Tetrachlorocuprate
Before and After Decolorization by Light
A
B
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-15-
1000 1200
WAVELENGTH np
Figure 17 Near Infrared Spectrum of Ethylammonium Tetrachlorocuprate
Before and After Decolorization by Light
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-16-
0.6 -
0.5 '
0.4
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800 ' 1000 ‘ 1200 
WAVELENGTH trp
Figure 18. Near Infrared Spectrum of Tetraethylammonium Tetrachloro-
cuprate Before and After Decolorization by Light
-17-
and they were not reproducible. (See Table IV page 41) Hence, the forma­
tion of the precipitate was attributed to reaction of the acetamidinium 
ion, and further study of (CH^C( N ^ ) 2] 2CuC*4 was abandoned.
The complete disappearance from the spectrum of the CuCl^ bands 
seemed to indicate that Cu(II) was being reduced to Cu(I). In accord 
with this, the following mechanism was proposed.
Initially, the action of light on the tetrachlorocuprate ion pro­
duces a chlorine radical and a Cu(I) species.
CuCl4 "' + hV ------> CuCl3"" + Cl-
The chlorine radical then abstracts a hydrogen from an acetonitrile 
solvent molecule, leaving an acetonitrile radical.
Cl- + CH3CN .....•* -CH2CN + HCl
The acetonitrile radical attacks another tetrachlorocuprate ion forming 
chloroacetonitrile.
-CH2CN + CuCl4 "" .....-» CuCl3"" + ClCH2CN
Several experiments were conducted in the hope of verifying this 
mechanism. Cuprous chloride dissolves readily in acetonitrile to give 
a colorless solution. Its spectrum (Figures 19 through 21) shows the 
same absence of bands in the near infrared and visible regions as those 
of the tetrachlorocuprate solutions after decolorization by light. If 
the photochemical reaction is run in a vessel having a volume larger 
than that of the solution so that an air space is present, total decolor­
ization does not occur. Instead, the solution acquires a greenish cast.
On the other hand, if the extra volume is filled with nitrogen rather
-18-
- 0.90.9 -
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0.7 - -0.7
- 0.6
- 0.50.5 -
_ 0.40.4 -
0.3-
-  0.20.2 -
-  0.10.1 -
Wavelength in m/j,
Figure 19. Ultraviolet Spectrum of Cuprous Chloride
-0.90.9 -
- 0.80.8 -
0.7 -
- 0.60.6 -
- 0.50.5 '
- 0.4
-  0.20.2 -
- 0.10.1 -
Wavelength in rrp
Figure 20. Visible Spectrum of Cuprous Chloride
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-19-
/
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0.1 -
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WAVELENGTH nyi
Near Infrared Spectrum of Cuprous Chloride
-20-
than with air, the reaction does proceed to a colorless state. Rebub- 
bling this solution with dried air, again, produces the green color.
The spectrum of the "nitrogen" solution compares with that of a solution 
decolorized with no air present. ThS spectra of the two green solutions 
are the same, but they differ from those of the "nitrogen" and "no air" 
solutions.
These results seem to support at least that part of the mechanism 
which calls for reduction of copper (II) to copper(I). It appears that 
some cuprous species is being photochemically produced, which, in the 
presence of air, is reoxidized to copper (II). Reoxidation might also 
be the explanation for the seeming reversal of the reaction on warming.
If the reaction proceeds as the mechanism suggests, a) a decrease 
in the number of free chloride ions should be observed, b) an increase, 
in the hydrogen ion concentration would be expected, and c) chloro- 
acetonitrile should be present in the solution. Unfortunately, none of 
the experiments undertaken to investigate these possibilities proved 
particularly successful.
Chloride ion determinations are hampered by the presence of ammonium 
species which interfere through complexation with standard silver nitrate 
methods. It was found, however, that qualitative results could be ob­
tained from turbidity measurements after precipitation of silver chlor­
ide with excess silver nitrate. From these, it can be said that the 
amount of free chloride does decrease as a result of the photochemical 
process, but no valid confirmation of the stoichiometry of the reaction
-21-
could be made. Free chloride values ranged from 2.5 per copper for the 
tetraethyl- and methylammonium salts to 3.0 for the ethylammonium com­
pound and could not be consistently reproduced. These are less than 
the value of 3.5 predicted from the mechanism. However, since all four 
chloride ions are initially free to react with silver nitrate, it appears 
that chloride ions are in some way being tied up as a result of the 
photochemical process.
Determinations of the pH of the tetrachlorocuprate solutions indi­
cate an increase in the hydrogen ion concentration after photochemical 
decolorization. The values obtained for the ethyl-and methylammonium 
salts agree with the expected ratio of one hydrogen per two coppers, 
but the ratio for the tetraethylammonium tetrachlorocuprate is more near­
ly one to four. However, the experimental errors are believed to be 
quite large, and no significance should be placed on the actual numer­
ical values.
Attention was next focused on the identification of chloroaceto- 
nitrile. Since the amount that would be formed according to the postu­
lated mechanism was less than the amount detectable by infrared or gas 
chromatography, the decolorized solutions were concentrated through 
removal of excess acetonitrile by distillation. A control solution of 
stoichiometric amounts of (X^H^^NCl, ClCHgCN, and CuCl in acetonitrile , 
was exposed to light and distilled in the same manner as the tetrachloro­
cuprate solutions. A comparison of the infrared spectrum of the control 
solution (Figure 22) with that of ClCHgCN and of (CgH^)^NCl in aceto-
-22-
nitrile, (Figures 23 and 24, respectively) indicates that chloroaceto- 
nitrile would not be destroyed under the conditions employed for reaction 
and identification of products. Hence, if chloroacetonitrile were formed, 
evidence for its formation should be obtainable from the infrared spec­
trum of the concentrated reaction mixture.
The infrared spectrum of a decolorized solution of tetraethylammo- 
nium tetrachlorocuprate (Figure 25) differs somewhat from that of a
solution that has not undergone photochemical reaction (Figure 26).
-I
Of primary importance is the disappearance of the bands at 752 cm and 
-I -i
665 cm , and the subsequent appearance of a band at 727 cm . A new
-I
peak also occurs at 1720 cm
-I
Although the 727 cm peak appearing in the infrared spectrum of
z
the decolorized solution very nearly matches the chloroacetonitrile peak 
-I
at 732 cm , positive identification of chloroacetonitrile as a product
-I
of the reaction cannot be made. The shoulder at 930 cm in the spec-
_x
trum of the decolorized solution might well correspond to the 927 cm
chloroacetonitrile band, but this would, then, represent a shift in the
-I
opposite direction to that observed in the 730 cm region which would 
be difficult to explain in terms of experimental error or solvent effects.
4 -i
There is also the question of the 1720 cm peak which is not present in
• the chloroacetonitrile spectrum.
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5000 3000 2000 1500 1300 1100 900 800 700 cm-1
Figure 22. Infrared Spectrum of ClCH2CN + CuCl + (C2H5)4NCl After Light
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C
E
N
T
 
T
R
A
N
S
M
I
T
T
A
N
C
E
5000
.gure 23. Infrared Spectrum of Chloroacetonitrile I
NJ
4>
I
1500 13003000 2000
Figure 24. Infrared Spectrum of Tetraethylammonium Chloride
P
E
R
C
E
N
T
 
T
R
A
N
S
M
I
T
T
A
N
C
E
 
2
 
P
E
R
C
E
N
T
 
T
R
A
N
S
M
I
T
T
A
N
C
E
5000 150,0 , 13,00 , IJfOO
25. Infrared Spectrum of Tetraethylammonium Tetrachlorocuprate Before Light
NJ
Ul
I
3090 j 2090 7p0 cm'
Figure 26. Infrared Spectrum of Tetraethylammonium Tetrachlorocuprate After Light
-26-
The second mechanism to be considered proceeds initially as does 
the first. Cu(II) is photochemically reduced to Cu(I), and a chlorine 
radical is produced.
CuCl4 "' +  hy .....-> CuCl3"" + Cl-
The chlorine radical, then, reacts with a second molecule of the tetra- 
chlorocuprate to form C ^  •
Cl- + CuCl4"" -----> Cl2 + CuCl3""
Chlorine could go on to combine with the solvent to give chloroaceto- 
nitrile.
Cl2 + CH3CN ...... ClCH2CN + HCl
This mechanism, like the firsts would predict an increase in hydrogen 
ion concentration and a decrease in the number of free chloride ions.
It should also be possible to detect chlorine or chloroacetonitrile in 
the solution.
Chlorine gas was found to dissolve readily in acetonitrile to give 
a pale yellow solution. The infrared spectrum of an acetonitrile solu­
tion of chlorine which had been exposed to light and concentrated by 
distillation (Figure 27) indicated that chloroacetonitrile could well 
be present as a product of this irradiation. The "chloro" band for
chloroacetonitrile occurs at 732 cm } while the band present in the
-I
spectrum of the irradiated chlorine solution falls at 728 cm . All 
other bands necessary for the identification of chloroacetonitrile occur 
at the correct positions.
-27-
It appears ,from the infrared spectrum that small amounts of chloro- 
acetonitrile may also be produced when a mixture of stoichiometric amounts 
of chlorine, cuprous chloride, and tetraethy!ammonium chloride in aceto­
nitrile is exposed to light. However, the infrared spectrum of this
solution after irradiation and concentration (Figure 28) shows an intense
-I -I
band at 1670 cm and a weaker one at 1230 cm , neither of which are 
present in the spectrum of the irradiated tetrachlorocuprate. Hence, 
it appears as if the actual photochemical process does not proceed exactly 
as the mechanism suggests.
An attempt to actually isolate the products of the tetrachloro­
cuprate reaction by fractional distillation resulted in the spectrum 
shown in Figure 29. If chloroacetonitrile had been present, an increase
O O
from 78 C . to 120 c. would have been expected. This, however, was not 
observed. Only one fraction was obtained which boiled at the same tem­
perature as acetonitrile. Also, the spectrum of this fraction does not
-I
show a band at 930 cm which should be present if chloroacetonitrile 
had been carried over with the solvent. Hence, it appears that chloro­
acetonitrile is not formed in. any appreciable amount as a result of the 
photochemical reaction.
P
E
R
C
E
N
T
 
T
R
A
N
S
M
I
T
T
A
N
C
E
 
?.
 
P
E
R
C
E
N
T
 
T
R
A
N
S
M
I
T
T
A
N
C
E
q s o p  ,isqo t Iioo20005000 , 3000
Figure 28. Infrared Spectrum of Cl^ + CuCl + (C^H^^NCl After Light
P
E
R
C
E
N
T
 
T
R
A
N
S
M
I
T
T
A
N
C
E
1500 1300 1100
_____i i i i________i
2000
I
ro
VD
I
Figure 29. Infrared Spectrum of Distillate from Tetrachlorocuprate Irradiation
-30-
It is well to consider, also, that since the equilibrium
2 CuCi4 "  ::= T 2 CuCi3"  + ci2
is believed to be present to some degree in solutions of the tetrachloro- 
cuprates, chlorine rather than the complex copper ion is the species 
which undergoes photochemical reaction. The visible spectrum of a 0.2 
molar solution of chlorine in acetonitrile (Figure 30) does show a weak 
absorption at 460 mfl, the region believed to be causing the photochem­
ical reaction. (The extinction coefficient at 460 nyA is 3.1 for the 
chlorine solution as opposed to 1070 for tetraethylammonium tetrachloro- 
cuprate at the same wavelength X This postulate can, however, be dis­
carded for essentially the same reasons given in the discussion of the 
previous mechanism.
0.9 -
0.8 -
0.7 -
0.6 -
0.5 -
0.3 -
0.2 -
0.1 -
WAVELENGTH m/t
Figure 30. Visible Spectrum of Chlorine in Acetonitrile
-31-
CONCLUSION
Of the three mechanisms proposed, the first appears to be the most 
nearly correct. With the use of more sophisticated methods, it might 
well be possible to isolate chloroacetonitrile. In any case, it can be 
said that alkylammonium tetrachlorocuprates undergo a photochemical re­
action when exposed to daylight or to light from a tungsten bulb. 
Copper(II) is reduced to copper (I) which is then present in the reaction 
mixture as some unidentified cuprous species. It is likely that chloride 
ion is the species which is oxidized; possibly to elementary chlorine 
which would be capable of reacting with the acetonitrile solvent.
It can also be said, in direct disagreement with the statement by 
Furlani, (See page 3 of the Introduction) that special precautions should, 
indeed, be taken during spectrophotometric measurements. In cases where 
exact values for the extinction coefficient are required, care should 
be taken to exclude daylight and artificial illumination during the prep­
aration of the solutions and the transfer of the sample to the instru­
ment. Absorbance values for a particular wavelength should be measured 
immediately after inserting the sample.
That such precautions are especially necessary in the near infrared 
region (the region of the tetrachlorocuprate d-d transition) of a Cary 14 
spectrophotometer is illustrated in Figure 31. During successive scans, 
the absorbance continued to drop until the band had virtually dis­
appeared. This rapid decrease in absorbtivity is not, however, evident 
during a scan of the visible region.
A
B
S
O
R
B
A
N
C
E
-32-
0.4"
0 . 2 -
1000 1100
WAVELENGTH mu
Figure 31. Decrease in Absorbance with Time
Note: The sharp peak at 1170 n>tz was shown to be an
instrument artifact and has been excluded from all 
but the initial scan.
-33-
Figure 32 shows a diagram of the optical system of a Cary 14 spec­
trophotometer. When1 operated in the visible region/ light from the 
tungsten lamp (c) is directed to the monochromator (F) and, then, alter­
nately through the sample cell (T) and the reference cell (Tt) to the 
photocell. Hence, only monochromatic light passes through the sample. 
When operated in the near infrared region, however, the light path is 
reversed. Radiation from the light source . (Y) is first divided between 
the sample and reference cells. It, then, reaches the monochromator 
where it is dispersed before being received by the lead sulfide cell. 
Consequently, in this region the entire spectrum of light, given off by 
the tungsten source is being focused on the sample cell, whereas, in the 
visible region, the light intensity is first considerably reduced by 
passing through the monochromator. It is important, then, to take into 
consideration the optics of the instrument being used as well as to 
make sure that the necessary precautions are taken to exclude light
from the sample.
PbS Cell
Tungsten
Lamp
■O —
■O ---
Sam.
Tungsten
Lamp
Figure 32. Optical Diagram of a Cary 14 Spectrophotometer^
-
34-
-35-
EXPERIMENTAL 
Preparation of Compounds
The tetrachlorocuprates were prepared, in this laboratory, by Dr. 
Emerson in the following manner. Stoichiometric amounts of the alkyl- 
ammonium chloride and anhydrous cupric chloride were separately dissolved 
in a minimum of ethanol. Upon mixing the two solutions, the correspond­
ing yellow alky!ammonium tetrachlorocuprate precipitated. These com­
pounds were recrystallized from ethanol-ether and dried in vacuo. When­
ever necessary, additional amounts of the compounds were prepared accord­
ing to this procedure.
Spectral Measurements
Spectra in the near infrared, visible, and ultraviolet regions were
measured on a Cary Model 14 recording spectrophotometer using 10.0 cm.
silica cells (Pyrocell Company), Beckman 1.0 cm. pyrex cells, and 1.0mm
quartz cells (Ultracell Company), respectively. Solutions of the tetra-
-3
chlorocuprates were approximately 10 molar in acetonitrile and were 
run against acetonitrile in each case. Infrared spectra were measured 
on a Beckman IR 5 spectrophotometer using matched sodium chloride liquid 
cells.
Decolorization of Tetrachlorocuprate Solutions with Light
Total decolonization of solutions of the tetrachlorocuprates in 
acetonitrile was effected through the use of 500 and 750 watt lamps in 
a Bausch and Lomb slide projector. The solutions, contained in a 10.0 cm. 
silica cell were placed directly in front of the projector, and a mirror
-36-
was placed directly behind the cell. Time for complete decolorization
-3
varied with the concentration of the solution, one half hour for 10 
molar solutions and approximately eight to ten hours for 10 ^ molar 
solutions.
Determination of Light Region Causing Decolonization
Colored filters were fitted over the end of the slide projector,
-3
and a 0.9 X 10 molar solution of acetamidinium tetrachlorocuprate was 
exposed to the filtered light. When a red filter. Corning 2-59, which 
transmits only light from 620 m/tto 3.7yt was used, no visible fading 
occured in two and one half hours. However, when the red filter was 
replaced by a blue-green filter. Corning 4-72, which transmits light 
from 340 to 600 mzt (A 460 m/0 and from 1.2 to 4.3/t, decolonization 
proceeded as before.
Determination of Change in Free Chloride and pH
Ten milliliters of 0.01 molar silver nitrate was added to five milli-
-3
liters of each of the following acetonitrile solutions: 0.83 X 10 M
C(C2H5)4NJ2CuCl4, 1.18 X IO-3M (C2H5NH3)2CuCl4, and 1.16 X 10"3M 
(CH3NH3)2CuCl4 . The resulting cloudy mixtures were, then, diluted to 
100 ml. with distilled water, and the turbidity (optical density) was 
measured at 400 ryz. with a Cary 14 spectrophotometer. Turbidity measure­
ments were also made with five milliliters of each of the solutions after 
decolonization by light. Using these'hbsorbance" values, the chloride 
concentrations were determined from a standardization curve of absorbance 
versus concentration for known solutions of tetraethy!ammonium chloride
-37-
(Figure 33). The results indicated a loss of free chloride for each of the 
tetrachlorocuprate solutions.
Measurements were also made of the pH of these solutions before and 
after light. In each case, three milliliters of the tetrachlorocuprate 
solution was diluted to 25 milliliters with boiled distilled water. 
Readings were taken ten minutes after the electrodes had been immersed.
In each case, an increase in the hydrogen ion concentration after the 
solution had been exposed to light was indicated.
Data for the chloride and pH determinations are given in Tables I 
and II, respectively.
Time Studies
A 10 ^molar solution of [(C2H5)^N]2CuCI^in acetonitrile contained 
in a 10.0 cm. silica cell was placed in the sample compartment of a 
Cary 14 spectrophotometer immediately after preparation. Successive 
scans were made from 1300 nyt to 800 tip*, over a period of one and one half 
hours during which time the absorbance at 1200 npa dropped from 0.51 to 
0.10. The time for one scan and return was approximately five minutes.
At no time was the solution removed from the instrument. Selected scans 
are plotted in Figure 31.
-4
In another experiment, the absorbance at 1200 m/6 of a 9.25 X 10 %
solution of acetamidinium tetrachlorocuprate was recorded at various 
times over a period of three days (Table III). After each reading, 
the sample was left in the instrument for ten minutes, and the change in 
the absorbance was recorded. Between readings, the sample remained in
A
b
s
o
r
b
a
n
c
e
Concentration of (C0Hs).NCl Moles/1 X 10
Figure 33. Standardization Curve for Determination of Chloride Ion Concentration
-39-
Table I. Determination of Free Chloride
A400 [cl J # of free Cl
____________ moles/1.__________ per Cu
(CH3NH3)2CuCl4 Before light 0.486 4.8 X 10-3
—3 -3 2.5
0.83 X 10 M After light 0.282 2.7 X 10
(C2H5NH3)2CuCl4 Before light 0.509 5.0 X
COIOi-H
— 3 -3 3.0
1.18 X 10 M After light 0.343 3.4 X 10
i* ^ —3
[(C2B5)4Nj2Cucl4 Before light 0.390 3.8 X 10
—3 -3 2.5
1.16 X 10 M After light 0.210 2.0 X 10
PH (,H+] due 
to HCl
(CH3NH3)2Cuci4 Before light 5.20 -3
0.24 X 10
0.83 X IO-3M After light 4.45
(C9i ghai1)9rerlf Before light 4.75 -3
-3
1.18 X 10 M After light 4.14
0.45 X 10
[(C2H5)4Nj2Cucl4 Before light 4.75 -3
1.16 X IO-3M After light 4.03
0.63 X 10
-40-
the cell, and no precautions were taken to exclude daylight or labora­
tory illumination. At the end of the three days, the absorbance at 1200 
m^tof a fresh sample of the original solution was taken, and the result
was compared with the previous measurements.
-4
In a third experiment, an 8.04 X 10 molar solution of acetamidin- 
ium tetrachlorocuprate was prepared in a blackened flask. The absorbance 
at 1100 m/< was measured immediately and again after twelve days. During 
this time, the absorbance decreased by 0.08 a.u.
Similar experiments with time were attempted in the visible region, 
as well, with the sample contained in a 1.0 cm. cell. Very little de- 
colorization was observed to take place in one to two hours.
Table III. Absorbance at 1200
Day Sample Time Intial
Absorbance 
After 10 min. in Cary 14
I I 11:30 am 0.720 0.675
1:30 pm 0.653 0.612
3:00 pm 0.595 0.560
4:30 pm 0.550 0.525
2 I 10:50 am 0.500 0.485
3 I 11:00 am 0.465 0.450
2 11:30 am 0.665 0.635
1:45 pm 0.615 0.595
3:00 pm 0.580 0.555
-41-
TABLE IV. d Spacings for Precipitate Obtained from Irradiation of 
Acetamidinium Tetrachlorocuprate
40 Sin 6 2 Sin 9 d = 1.5418/2 Sin 6
51.5 0.2228 0.4456 3.46
53.7 0.2321 0.4642 3.32
70.5 0.3028 0.6056 2.55
86.6 0.3689 0.7378 2.09
105.5 0.4442 0.8884 1.74
115.1 0.4814 0.9628 1.60
-42-
Literatnrp OitRd
,
1. F . Basolo and R. G . Pearson3 Mechanisms of Inorganic Reactions, John 
Wiley and Sons3 Inc. 3 New York3 1967, pp. 662-663.
2. J . Cislova3 J . Gazo3 and G . Ondrejovic3 Z . Chem., _6, 429 (1966).
3. F . Cotton and G . Wilkinson, Advanced Inorganic Chemistry, 2nd Edition3 
Interscience Publishers, New York3 1966, p. 899-900.
4. G. Ho Faye3 Can. J. Chem., 44, 1643 (1966).
5. C. Furlani and G. Morpurgo3 Theoret. Chim. Acta, (Berl.) I, 102 (1963) .
6. L. Helmholz and R. F . Kruh3 J. Am. Chem. Soc., 74, 1176 (1952).
7. W. Schneider, Proceedings Seventh International Conference on 
Coordination Chemistry3 Abstract 1B7, (1962).
8. N. V. Sidgewick3 The Chemical Elements and Their Compounds, Oxford 
University Press3 London3 1950, pp. 161-162.
9« R. D. Willett, 0. L. Liles, Jr., and C . Michelson, Inorganic Chemistry, 
6, 1885 (1967).
10. R„ D. Willett, J . Chem. Phys., 41, 2243 (1964).
Instruction Manual, Cary 14 Spectrophotometer.11.
N378
M316 M annlla, Kathleen D. 
cop.2 A photochemical reactioi
of alkylammoniunt tetra- 
chlorocuprates
u moq mu ­  AODnKfttt
N 31« 
M  S i t ?
C O r 3 ~