THE SYSTEM PYRIDINE-HYDROGEN CHLORIDE AS AN ACID MEDIUM BY HYMAN MITCHNER
A THESIS.SUBMITTED IN PARTIAL FULFILMENT OF THE REQUIREMENTS FOR THE DEGREE OF MASTER OF SCIENCE in the Department of CHEMISTRY
We accept this thesis as conforming to the standard required from candidates for the degree of MASTER OF SCIENCE
Members of the Department of Chemistry
THE UNIVERSITY OF BRITISH COLUMBIA
APRIL, 1953 ABSTRACT
Pyridine salts were investigated as acids in the pyridine system. The mono and the dihydrochloride salts were found to be the best dissolving reagents for the a metals and the sulphides used. Pyridine hydrochloride was most effective in the molten state, whereas pyridine dihydrochloride was found to be quite reactive at room temperature when dissolved in a chloroform solution.
Side reactions were investigated and were found to occur only with Mg, Al, and Zn with molten pyridinium chloride.
The complex salts, (C5HeN.H)8[MnClB3c5HBN and
(C5HBN.H)HgCl4 were isolated and investigated. 1
I
AKNOWLEDG-EMENT
Sincere appreciation is expressed to Dr. K. Starke for his'encouragement and help in making this work possible. TABLE OF CONTENTS
PAGE
INTRODUCTION
1. Water as an unique solvent 1
2. Historical approach to reactions in nonaqueous
systems
a. . Confusion and misconcepts in the acid-
has e theory...... 1
b. Bronsted theory of acids and bases.. «,...© 4-
3. Disadvantages of water as a solvent...... 5
4. Alms of investigation ...... 5
EXPERIMENTAL
1. Purity of reagents 7
2. Preparation of pyridinium salts
a. Pyridinium nitrate 7
b. Pyridinium trlchloroacetate 7
Co Pyridinium thiocyanate 8
d. Pyridinium oxalate 8
e. Pyridinium fluoride...... 9
f. Pyridinium chloride 9
g. Pyridine dihydrochloride...... 11
h0 Other salts of pyridine...... 11
3. Methods of using C5HBN.HC1 and CBHBN.2HC1
a. Molten C5H5N.HC1 and molten CBH6N.2HC1... 12 -11
b. Room temperature systems
1. C.6HBN.HCl in pyridine,,...... 14
lie Saturated chloroform solutions of
CBHBN.HC1 and CBHBN.2HC1...... 15
c. The reaction of metals in molten CeHeN.HCl
and the reaction of metals with a satur•
ated chloroform solution of C6HBN.2HC1
at room temperature..... 16
Solubility of C5HBN.HC1 and CBHBN.2HC1 in. - - chloroform at 23°C 17
Absorbency - wavelength, investigation of the. ,
chloroform solutions of GBHBN0HC1 and
CBH5N.2HC1 17
Side reactions
a. Materials
i. Preparation of 4-pyridyl-pyridinium
dichloride 20
ii. Purification of synthetic quinoline. 20
iii. Preparation of quinoline hydrochlo•
ride 21
b. Product of side reactions
1. Isolation...... 22
ii. Examination for bipyridyls. 23
ill. Decomposition product of 4-pyridyl-
pyridinium dichloride in NaOH...... 24 -iii*
iv. Further reactions of Zn + C5HBN.HC1
reaction product...... •« 24
c. Reaction of quinoline hydrochloride with
metals • . ,.
i. Side reaction product ofquinoline,
'hydrochloride plus zinc 25
ii. Reaction of quinoline hydrochloride
product with acetic. anhydride 25
7. Reactions with possible analytical application
a. Materials - preparation.of. Mn.(OAc )3...... 26
b. Reaction of CBHBN.2HC1 in CHC13 with MnS
i. Nature of sample of, MnS 26
ii. Reaction of. manganese, salts,. metal. •.
and Mn08 with a saturated chloro•
form solution of CBHBN.2HC1..... 27
iii. Heating of a fresh sample of MnS..... 27
iv. Effect of oxidizing agents ..... 27
v. CBHBN.2HC1 in CHC1S on Mh(OAc)8."..... 28
vi. Isolation of the substance confer—
ring.the green colour to the
G5H5N.2HCI in GHCI3...... 28
vii. Analysis pf the green crystals...... 29
viii. Limit of visible colour...,..,,...... ,. 29
ix. Colourometric analysis..,,.,....,... 29
x. Variation of colour with time.33 4
-iv-
c. Reaction of CBHeN.2HCl in CHC13 with HgS.
i. Isolation of the•mercury salt...... 33
ii. Analysis of the mercury salt...... 33
III DISCUSSION
lo Pyridine salts as rection media
a. Type of reaction most desirable. 37
b. Examination of the pyridine salts pre•
pared 37
2. Methods of using CBHBN.HC1 and CBHBN.2HC1
a. Molten state 38
b. Room temperature systems 39
3. The nature of CBHBN.2HC1 in CHC13 Al
4. Side reactions a. Possibility of formation of bipyridyls... 42
b. Ring cleavage...... 43 c. Limitations due to side reactions 45
5. . Reactions with possible analytical application
a. Trivalent manganese
i. Formulation of the compound
(C5HBN.H)sCMnClB3CBH5N...... 46
ii. Explanation of reactions of manganese
sstXts© •»••••«*««»••«•««••»• • • •«o«• © • *
iii. Determination of trivalent manganese
by colourometric analysis of the
green complex... 48 -V-
b. The mercury-pyridine complex
.1. Formulation of the compound
(CBHBN»H)gHgCl^••.••••.«•«..««*»•«.. 50
ii. Application of solubility of HgS.... 51
6. Conclusions .. • ... 51
XV BIBLIOGRAPHY. «....-...... 0...... 0...... • « 52
TABLES ......
Table I Reactivity of sulphides in molten C5HeN.HCl
and molten CBHBN.2HC1...... *.....*...••..... 13
Table II Reactivity of sulphides in CBHBN.HCi in
pyridine. • 14-
Table III Reaction of sulphides in saturated chloroform
solutions of CBHBN.HC1 and C5HBN.2HC1 15
Table IV The reaction of metals In molten CBHeN.HCl
and the reaction of metals with a saturated
chloroform solution of C5H6N.2HC1 at room
temperature • 16
Table V Reaction of manganese salts, metal and Mn08
with a saturated chloroform solution of
CBH BN.2HC1...... 27
0 -vi-
FIGURES
Figure I Absorbency-wavelength plot for a saturated
chloroform solution of C6HBN.2HC1 and a
diluted chloroform solution of CBHBN.HC1... f'lQ i
Figure II AbsorbencyTwavelength.plot for a diluted •
chloroform solution of C5HBN.2HC1 .. 19
Figure III Spectral-Transmittance curve of green tri•
valent manganese complex la a saturated
chloroform solution of CBHBN.2HC1.. 31
Figure IV Concentration-Transmittance curve of green
trivalent manganese complex in 10 cc. of a
saturated chloroform solution of C5H5N.2HCI.
Trivalent manganese obtained by reduction
Figure V Concentration-Transmittance curve of green
trivalent manganese complex in 10 cc. of a
saturated chloroform solution of CBH5N.2HC1.
Trivalent manganese obtained by oxidation
of divalent > manganese, os »».»»•«.»..... •«... 34
Figure VI Variation of colour of green trivalent man•
ganese complex with time. .35 -1-
I - INTRODUCTION
1. Water as aaunique solvent.
Of all our known solvents, the one most used is
water. As a solvent, water is considered to be unique.
Its physical properties, such as its capacity as a gen-
eral solvent for salts and its power of electrolytic
dissociation, its low molecular elevation constant, its
high boiling point, and its heat of fusion, heat of
volatilization, critical, temperature, specific heat,
association constant, and dielectric constant with values
so much higher than the corresponding values for other
substances, all tend to remove it far from other sol•
vents and to place it in a class by itself. Furthermore,
because of its abundance, it is one of the most studied
substances known. ... . •
2. .Historical approach to reactions in nonaqueous systems.
a. Confusion and misconcepts in the acid-base
theory.
Following the discovery of the voltaic cell by
Volta early in the 19th century, considerable work was
done to differentiate between those substances which
conduct an electric current in aqueous solutions and
those which do not. Electrochemists of the middle 19th
century such as Hittorf and Kohlrausch studied the -2-
behavlour of the so-called electrolytes such as hydrogen
chloride and ammonia in their liquid states. When these
latter two substances were found to be nonconductors,
it was promptly assumed that they did not break up into
ions and that only in water did they possess character•
istic chemical and electrochemical properties, and hence
only in water was it possible to effect ionic reactions.
The criterion for the pure substance to be able to.act
as an acid per se was thus related to its ability to
conduct electricity.
The role of water as an ionizing medium In electro•
lytic dissociation intrigued many investigators.
Arrhenius (1884) attempted to completely explain electro•
lytic dissociation by placing the emphasis on the role 37 of water. Werner , the originator of the complex compound theory, also tried to explain what happened when a nonelectrolyte dissolved in water to give a conducting solution. He introduced the terms "anhydro base" and "aquo base", as well as "anhydro acid" and "aquo acid". "Anhydro bases (such as ammonia) are compounds which combine with the hydrogen ions of water in aqueous solu• tion, and thereby cause a shift in dissociation equili• brium of water until their corresponding characteristic hydroxyl ion concentration has been reached." "Aquo bases (such as potassium hydroxide) are addition compounds of water, which in turn dissociate in aqueous solution to give hydroxy! ions." Anhydro acids (such as HCl) are "compounds which in aqueous solution so combine with hydroxyl ions of water as to cause a shift in dissocia• tion equilibrium of the solvent water until the corres• ponding characteristic hydrogen ion concentration has been reached."
It was not long, however, before the limitations of the Arrhenius concept were recognized. Work on non- 36 aqueeus solutions by Walden in Europe and contemporary is activity in the United States by Cady, Franklin , and
Kraus made a revision of acid and base theory necessary.
It was shown that numerous solvents yielded conducting
solutions and that all substances exhibit^solvent prop• erty to a greater or lesser degree. Some compounds regarded as salts in the aquo system were found to behave as acids or bases in nonaqueous systems and therefore, certain changes with respect to the definitions of classes of compounds in terms of the solvent employed, were obviously necessary.
Ammonium salts behave as ammono-acids in liquid ammonia as the solvent , whereas acetates behave as 9 aceto-bases in glacial acetic acid . In general, "onium"
salts exhibit acid characteristics if the corresponding anhydro base is the solvent. b. Br^nsted theory of acids and bases.
One of the most useful theories proposed to coordinate these more advanced investigations was one proposed by Brpnsted . The phenomenon of acidity was simply considered as a matter of competition between the solvent and the acid anion as bases for the proton:
A^=±B + H+
Thus an acid would be any molecule or ion which acted as a proton donor, a base, any molecule or ion which acted as a proton acceptor. Acidity in basic solvents would be due to the formation of the corresponding "onlum" ion. In ammonia there is formed the ammoniated hydrogen ion or ammonium ion as the bearer of acidity:
+ + NH3. + H NH4
Similarly in aniline there is formed the anllinium ion
+ (C6H6NHS«H } and in pyridine the pyridinium ion
+ (CBH5N.H ).*
The natural experimental extension was to investi• gate the solid and fused states of the "onium" salts in which the "onium" ion must exist, to see if the salts of the solvent exhibited therein the properties of acids.
Investigations carried out with fused pyridinium chlorfede
* The Brc^nsted theory, of course, is limited because it takes into consideration only protonic subs• tances, but where applicable it is extremely useful. As the solvents considered in the investigation herein are protonic, the Brj^nsted theory is used.
i demonstrates conclusively that an extension to the molten state of the "onium" reactionsu&s perfectly justi• fied. Furthermore, all other "onium" salts Investigated, have been found to behave in a similar manner .
Disadvantages of water as a solvent.
The high dissociating power of water can prevent the formation of neutral molecules or complex ions which might form in nonaqueous reactions, but dissociate in the presence of water. Undesirable dissociation may even be followed by a reaction with the solvent as a result of water itself being always dissociated. Thus, in order to obtain an overall picture of reactions in nonaqueous solvents a series of reactions in which water was not formed was carried out.
Alms of investigation.
a. It was proposed to investigate the acid salts of pyridine as reaction media for metals and common sulphides.
b. It was proposed to investigate further those salts which showed, in the molten state, good reactivity for the metals and sulphides used, with the aim of obtain• ing a solvent system which could be used at room temper• ature.
c. It was proposed to investigate any side reactions which might limit the applicability of the reaction media used. d. It was proposed to investigate any reactions that might suggest possible analytical applications. II - EXPERIMENTAL
1. Purity of reagenta.
The reagents used In all preparations were of
"Reagent Grade" quality* The pyridine was further pur•
ified by refluxlng for two hours over sodium hydroxide
pellets to remove any water that might have been absorbed
from the air. The pyridine was then distilled and the
fraction coming over between 115-ll6°C. collected.
2. Preparation of pyridinium salts. 1*33
a. Pyridinium nitrate (C5HBN.HN03)
Slightly more than an equivalent amount of concen•
trated nitric acid was added to 100 gm. of pyridine and
' the resultant mixture evaporated on a steam bath till
viscous. Upon cooling, a crystalline product was obtained
which was treated with ether, filtered by suction and
washed again with ether. The pyridinium nitrate was then
recrystalllzed from absolute alcohol and white needles
of large size were obtained, m.p. 115-116°C. The yield
was quantitative. No melting point has been previously
reported.
b. Pyridinium trichloroacetate (C6HeN.GClaC00H)
To prepare pyridinium trichloroacetate, equivalent
amounts of pyridine and trichloroacetic acid were used.
Particular care was taken to keep the trichloroacetic —8** acid free from moisture. The reaction was carried out in an ice bath. The pyridine was added slowly with stir• ring to the cooled trichloroacetic acid. A vigorous reaction ensued with the production of a yellow-brown semi-solid. After the reaction was complete, the product was heated on a steam bath till liquid and cooled again in the ice bath, where it formed a yellow slush. The cold residue was treated with absolute ether, filtered and washed twice with additional amounts of absolute ether. The buff coloured residue was recrystallized from-absolute alcohol and gave white, planar (mica-like) crystals which decomposed on heating at 111-112 C. Reitzenstein report• ed a melting point at 112°C.
13
c. Pyridinium thiocyanate (CeH6N.HSCN)
Pyridinium thiocyanate was prepared by the add• ition of alcoholic solutions of equivalent amounts of pyridine hydrochloride and ammonium thiocyanate. On addition of the two alcoholic solutions, NE^Cl precip• itated out. This was filtered and the product recovered by evaporating the filtrate-to a small volume. The crystals thus obtained were filtered and washed with ether.
They were easily recrystallized from ethanol to yield white planar crystals of melting point 98°C. No melting point has been previously reported. 3 8 > 3 X
d. Pyridinium oxalate ((C5H5N)sC2Hs04)
Pyridinium oxalate was prepared by the addition of pyridine to an acetone solution of axalic acid. A bulky
white precipitate was obtained which was filtered, washed
with acetone and recrystallized from absolute alcohol to 3 i
give white crystals of m.p. 152-153°C. Pfeiffer reported a. melting point at 153'vC.
e. Pyridinium fluoride
Pyridinium fluoride was not prepared. An- attempt
to prepare it along the line of pyridinium nitrate was not successful. Equivalent amounts of pyridine and hydrogen fluoride (in the form of 48$ HF) were mixed
in a stainless steel beaker and the resulting solution
evaporated. No product was obtained. A second attempt
to prepare the hydrofluoride was made by adding pyridine
to aqueous HF, but this time acetic anhydride was added
to reacg with the water. The entire mixture was extracted
with ether to remove the acetic acid and to leave behind
the pyridinium fluoride, which, if it behaved as other pyridine salts, would be insoluble in ether. No product was obtained. The final method tried Included the above
two procedures, but an excess of hydrofluoric acid was used to the extent of six moles of HF to one mole of pyridine. No pyridinium fluoride could be isolated.
S 3 $ 3 mm
f. Pyridinium chloride (CBHeN.HCl)
Pyridinium chloride is hygroscopic and it was found best to prepare it Just before use. Three methods were found to work quite readily! -10-
i. Using a Kipp generator, HCl gas .was evolved by reacting concentrated sulphuric acid with ammonium chloride. The gas was dried by passing it through two gas washing bottles containing concentrated sulphuric acid and was then introduced into ,a solution of pyridine in dry ether. Pyridinium chloride precipitated almost immediately as a white salt*. It was filtered in a
Buchner funnel and stored in a vacuum deslcaator over anhydrous magnesium .perchlorate.
Ii. Concentrated hydrochloric acid was added to a slight excess of pyridine and the resulting aqueous solution of pyridinium chloride was distilled. The dis• tillate coming over in the range 218-218.5°C was collect•
ed as CBH6N.HC1. The solidified salt was then dissolved in absolute.alcohol and crystallized by cooling. The salt remaining in the mother liquor could be precipit• ated with dry ether.
iii. Anhydrous ether was saturated with dry
HCl and placed in a separatory funnel. The solution was slowly added to freshly distilled pyridine until the pre• cipitation of the pyridinium chloride was almost complete, whereupon this solution was treated as in (ii). The
* Pyridine should be present in excess to prevent the possible formation of C6H5N.2HC1. -11- melting point of pyridinium chloride prepared by any of these three methods was found to be l43-l44°C. This value was identical to that reported by Audrieth and a Long o 16
g. Pyridine dlhvdrochloride (CBHBN..2HC1) Dried HCI gas from a Kipp generator was passed through freshly distilled pyridine contained in a three- neck flask fitted with both a stirrer and a reflux condenser. The pyridine hydrochloride formed first and the solution, was then kefit at a temperature sufficient to maintain a homogeneous solution. HCI was passed through continuously and, after the formation.of the
CBHBN.HC1 was completed, the temperature,of the flask was lowered.so that it was maintained Just above that point at which crystals started to form. This was cont• inued till the temperature of the solution in the flask fell to about 48°C. On cooling a crystalline mass solid• ified, m.p. 46-47°C. and decomposition point 55°C. These 16 values agreed closely to those of Kaufler and Kunz m.p, 46.7°C. and decomposition point 55°C. By direct weighing it was observed that the pyridine had taken on two moles of HCI per mole of pyridine. The crystals can be reprecipitated from anhydrous ether and alcohol as an oil and white needle-like crystals. h. Other "Se4-d- salts of pyridine
8 9,36 i. HsFe(CN)6 and H4Fe(CN)6; * -12- 13
ii. C4H40e (furoic acid)
39 iii. G8H604 + 2CeHBN (phthalic acid)
iv. H3P04 3 3 S3
v. HsMo04 16 >35*3 9
vi. HCr06 1»S3
vii. HsS04 3 0 viii. HI 31 ix. HBr
Methods of using C6HSN..HC1 and CBH6N.2HC1.
a. Molten CBHBM.HCl and molten CBH5N.2HC1
A series of test tube reactions we>$^ carried out using the reagents mentioned. The reactivity of these reagents on metal sulphides wf|®$. investigated with the
evolution of HSS as the criterion of reaction. Lead acetate paper was used to check the evolution of the
HaS gas. TABLE I
Reactivity of sulphides in molten CKHKN.HC1 and molten
CBHeN.02HCl
Sulphide CBHeN.HCl at 175°C CBHBNi.2HCl at 53°C.
Colour of Reactivity Colour of Reactivity solution solution
MnS 1 2 ZnS 1 1 PeS orange 1 yellow 1 CdS 1 1 CoS blue-green 1 blue-green 1 NiS blue 1 blue 1
SnS 1 i
4 SnS8 4
PbS 1 3
SbsS3 1 1
BisSa 2 3
As8S3 3 4 CuS It. yellow 1 It. yellow 2
Ag8S 2 2 HgS 1 1
1 - quite reactive: 2 - reactive: 3* slightly reactive: 4 - unreactive. b. Room temperature systems
i. CKHKI.HCX In pyridine (6N in GBH5N.HC1)
TABLE II
Reactivity of sulphides in CBHBN.HC1 in pyridine
Sulphide Colour of Reactivity solution
MnS 3
ZnS 3
FeS yellow 2
CdS . 3
GoS blue-green 2
N1S ... 4
SnS 3
SnS8 4
PbS 3
Sb8Sa 4
BiaS3 3
As3S3 4
CuS 3
AgsS 4
HgS 4
1 - quite reactive: 2 - reactive: 3 - slightly reactive: 4 ~ unreactive* -15- 11o Saturated chloroform solutions of
C8HSN.HC1 and CBHBN.2HC1
TABLE III
Reaction of sulrjhides in saturated chloroform solutions of
C5HBN..HC1 and C5HBN.2HC1
Sulphide CBH6N.HC1 in CHC1S CBHBN.2HC1 in CHCla,
Colour of Reactivity Colour of Reactivity solution solution MnS 3 dark green* 1 ZnS • 1 1 FeS yellow 2 yellow 1 CdS 1 1 CoS blue-green 2 blue-green 2 NiS blue 2 blue 1
SnS 3 1
4 SnSa 4
PbS 3 1
SbaSa 2 1
BigSg 4 3
AssS3 4 3 CuS It. yellow 3 It. yellow 2
AgsS 3 3
HgS 2 1 1 «• quite reactive:: 2 - reactive: 3 - slightly reactive: 4 - unreactive. * A freshly prepared sample of MnS gave a colour• less solution. An old sample gave a dark green solution. c. TABLE IV
3 > 3 » 36.
The reaction of metals in molten 0kHKN..HC1 and the reaction of metals with a saturated chloroform solution
of CBH/BN.2HC1 at room temperature
Metal Molten CH«N«HC1 CKHBN,2HC1 In CHCla (25°C;) Colour of Reactivity Colour of Reactivity solution solution Al brown 1 2 Bi 4 3 Mn 1 • 1 Qd yellow 1 yellow 1
Co blue-green 2 blue-green 2 Cr pink 2 pink 2 Cu yellow 1 yellow 1 Mg brown 1 1 Hg 4 4 Ni blue 2 blue 3 Pb 2 3
Sb 3 3 Sn 2 2 Zn brown 1 1
1 - quite reactive: 2 - reactive: 3 - slightly reactive: 4 - unreactive. -1T-.
After driving off the chloroform, the only changes obs•
erved in the reactions carried out with the CBHBN.2HC1
in CHG13 on metals, were that the solutions of Al, Zn, and Mg turned brown. This was the same colour as was
observed with the molten C5HBN.HC1 on the same metals.
Solubility of C^HBN.HC1 and CfiHEN.2HCl In chloroform at
23° C.
The solubilities of pyridinium chloride and pyridine dihydrochloride were measured by saturating a chloroform solution of known volume with these reagents and pre•
cipitating the pyridine content as 2C5HBN.CuCls, using
excess CuCl8 in a water solution. The chloroform was driven off by heating on a steam bath. The values determined were:
CBHBN.HC1 - 123 gm. per 100 cc. of chloroform
CBHBN.2HC1 - 108 gm. per 100 cc. of chloroform
Abserbencv - wavelength investigation of the chloroform
solutions of C5HEM.HC1 and CBH5M.2HC1. o The wavelength region \ = 3300 - 5300 A was inves• tigated on a Beckmann Spectrophotometer. The absorbency was taken as the dependent variable. A pure chloroform solution was used as the standard. Figure I represents
a saturated solution of CBHeN.2HCl in CHC13 and a diluted
chloroform solution of CBHBN.HCl, Figure II represents a dilution of the saturated chloroform solution of
CBHBN.2HC1. I L—-„ i ,_J C_. i _Ji L.. I Ll. L I j i L_ 1 ! 1 3*00 HObO' .. HS0O • j'oOO
WAVELENGTH X ^° V
FIG.UR€ Io' ABSORBENCY-WAVELENGTH PLOT FOR A SATURATED CHLOROFORM SOLUTION OF Cc;H^N«2HCl
• AND A DILUTED CHLOFORM SOLUTION OF C5H5N0HCI0 • • FIS.URE II, ABSORBSNCY-WAVELENGEH "PLOT FOR A DILUTED CHLOROFORM SOLUTION OF C^H^N02HClo Side reactions.
a. Materials
i. Preparation of 4-pyrldyl-pyrldinlum 17/
dichloride
One hudred and fifty grams of thionyl chloride were slowly added to 50 gm. of pure pyridine. The mixture was kept cool in an ice bath. The resultant yellow solution was refluxed for five hours, whereupon it grad• ually assumed a dark brown colour. The brown solution was distilled under vacuum; the temperature was slowly raised to 100°C. and was maintained there for one hour.
The dark brown residue in the distilling flask was stir• red up with 50 cc. of absolute alcohol and filtered by suction. There remained a buff coloured granular residue which was dissolved in dilute hydrochloric acid, boiled and filtered. The filtrate was evaporated till crystal• lization just commenced. At this point alcohol was added to the cooled mixture. There resulted a faintly yellowish mass of crystals of 4-pyridyl*pyridinium dichloride which were filtered and washed with alcohol. 38
ii. Purification of synthetic quinoline
One hundred cc. of quinoline were dissolved in
1200 ml. of dilute hydrochloric acid and heated to 60°C,
whereupon a solution of 140 gm. of ZnCl8 in 24o ml. of dilute hydrochloric acid was added. -21-
2C9H7N + ZnCl8 + 2HC1 — —* C(09H7N)8ZnCu3Ha
A white precipitate of the quinoline chlorozincate soon began to form and the well stirred mixture was cooled in an ice bath. The crystals were separated by filtering with suction and washing with dilute hydrochloric acid.
The white crystalline precipitate was transferred to a beaker and 10% NaOH solution added till the precipitate
of Zn(OH)8 dissolved. The solution was then extracted , with six 100 ml. portions of ether and the combined ether
extracts were dried with 20 gm. of anhydrous magnesium
sulphate. The ether was distilled off. The water con• denser was replaced by an air condenser and the fraction boiling between 237-238°C. was collected. Redistillation
in vacuo gave a clear colourless distillate of quinoline.
iii. Preparation of quinoline hydrochloride
(0.9H7N,HC1)
Fifty cc. of the purified quinoline were dissolved
in 100 cc. of ether and dry HCI passed through the solu•
tion to precipitate the quinoline hydrochloride which
is difficultly soluble in warm ether and soluble in hot 10
ether . The precipitate in ether was cooled in an ice
bath and filtered, using suction. The precipitate was
washed twice with cold ether and dried In a vacuum des-; ,
locator, m.p. 134°C. The same value for the melting point was obtained by Erkstein 10 . -22-
b. Product of side reactions
i. Isolation
The reactions of Zn, Mg, and Al with molten pyridinium chloride result in a brown solution. The reaction with Zn has been studied as typifying those
resulting in side reactions. Pyridine plus ZnCl8, pyridine dihydrochloride plus Zn, pyridine hydrochloride
plus ZnCl8 did not result in a solution of brown colour.
The brown colour was typical only of the reaction of the metal with pyridine hydrochloride.
Thlrty^five grams of CBHeN.HCl were reacted with
an excess of zinc dust at 175°C8 As zinc dust was added, the solution turned light-yellow, but rapidly turned brown on the addition of more Zn. On cooling, the mix• ture was treated with 600 ml. of 6N. NaOH. A brown oil
separated on the surface of the solution0 The oil had a very strong odour of pyridine. It was separated from the alkaline solution by a separatory funnel. The brown oil was washed with five 50 ml. portions of 6N
NaOH by shaking in the separatory funnel and drawing off the NaOH. The pyridine in the brown oil was removed by evaporating the solution on a steam bath till a tarry brown sbstance remained. This product did not crystallize from any of the common organic solvents.
It was found that the presence of pyridine with the residue Interferred with the isolation of the reaction -23- product and it was necessary to remove it completely before further purification could be made, Repreclpitation in an amorphous form could be accomplished if the the tarry residue from which the pyridine had been completely evap• orated was treated as follows:
1, The residue was dissolved in chloroform and the solution evaporated on a steam bath till it became a sticky mass,
2, The residue was taken up with about six times its volume of carbon tetrachloride, at which point some buff?-coloured precipitate appeared.
3, Low boiling petroleum ether (30-60°C,) was added to complete the precipitation.
The maximum amount of buff.coloured residue obtain• ed was about 6% of the starting weight of the pyridinium chloride. The powder obtained by the described method did not melt up to 350°C, It was found to be soluble in aqueous acids, ethanol, and chloroform giving a red coloured solution in each, but insoluble in benzene, water and ether,
iia Examination for bipyridyls
The sodium hydroxide solution, the brown oil, and the buff coloured residue were extracted with ether for bipyridyls (all of which are soluble in ether). On evaporation of the ether, no bipyridyls were obtained. -24-:
Iii. Decomposition product of 4-pyrldyl-
pyrldinlum dlchloride in NaOH
4-Pyridyl-pyridinlum dlchloride was added to a 6N
solution of sodium hydroxide. An intense yellow colour
appeared and, on the addition of more 4-pyridyl-pyridinium
dlchloride, a red-brown precipitate formed,, An aldehydic
cinnamon odour was noticeable. The precipitate dissolved
in acetic acid to gfeve a red-brown colour, very much like
ythat of-the buff coloured reaction product from pyridine
hydrochloride and zinc. However, this new product formed
a derivative with 2-4-dinitrophenylhydrazine, whereas
the pyridinium chloride reaction product with zinc did
not form any precipitate with this latter reagent.
iv. Further reactions of Zn + CBH6N.HC1
reaction product
A solution of the reaction product in 6N acetic
acid gave the following reactions:
1. Decolourized a 2% KMn04 solution.
2. Became colourless when allowed to stand after
treating with 30$ H808.
3. Gave a yellow-brown precipitate with 20%
Br8 in aqueous KBr.
40 Gave a dark brown precipitate with 20$ I3
in aqueous KI.
On isolation of the bromine derivative as a brown
powder, no melting point could be determined. Analysis showed 5.29$ nitrogen.
c. Reaction of quinoline hydrochloride with metals
i. Side reaction product of quinoline
hydrochloride plus zinc
Quinoline hydrochloride was found to react with all metals.above hydrogen in the Electromotive Series to give a dark orange to red solution. The reaction of quinoline hydrochloride and zinc resulted in a dark red solution which, when poured into water, formed a dark red residue. This residue was powdered In a morgar and pestle and washed with cold water. It was dissolved in concentrated hydrochloric acid solution and reprecip- itated with chloroform. The residue was again washed with cold water and allowed to dry in a vacuum desic• cator. On standing, some of the substance appeared to turn tarry, but on powdering the residue, an apparently uniform dark' red substance was obtained, m.p. 126°C.
The yield was 60% of the weight of the quinoline hydro• chloride reacted.
ii. Reaction of quinoline hydrochloride
product with acetic anhydride
The red product obtained from the reaction of
quinoline hydrochloride and zinc was treated with an
excess of acetic anhydride and refluxed for two hours.
Water was added to decompose the excess acetic anhydride
and the solution was neutralized with dilute ammonium -26- hydroxide. A yellow-brown colloidal precipitate resulted which was centrifuged, washed with very dilute ammonia, filtered and dried. This substance did not melt up to
350°C. At this point research into these side reaction products had to be abandoned because of the lack of time.
Reactions with possible analytical application.
a. Materials - preparation of Mn(OAc)3
One hundred cc. of acetic acid were heated to boil• ing in an evaporating dish and 6.9 gm. of anhydrous
Mn(0Ac)8 added. After the acetate had been dissolved,
1.6 gm. of KMn04 was added slowly with stirring. The solution turned dark brown and after heating for a few minutes was allowed to cool. The solution was concen• trated and on cooling a crop of manganiacetate crystals was obtained. The first crop was washed with dilute acetic acid and recrystallized twice from glacial acetic acid. The crystals were dried in a vacuum desiccator over KOH,
KMn04 + 4Mn(0Ac)s + 8H0Ac — <+
5Mn(0Ac)3 + KOAc + 4H30 ..
b. Reaction of C6HBN,2HG1 in CHC13 with MnS
i. Mature of sample of MnS
Freshly prepared MnS reacted with the reagent used, but did not colour the solution. A sample of MnS pre- pared at least one year previously, reacted to give an intense green solution. . ii. TABLE V
Reaction of manganese salts, metal and MnOa with' a sat•
urated chloroform solution of CBHBN.2HC1
Substance Reaction Colour of solution
MnC03 Very reactive Dark green
MnS04 No visible reaction
Mn(N03)a Some reaction Dark green
MnCl3 No visible reaction
Mn03 Very reactive with Dark green
evolution of Cls
Mn(OAc)a Some reaction
Mn Very reactive
The salts, manganese metal, and MnOs were also tested
with CBHBN.HC1 in CHC13 but did not give the green colour.
\ iii. Heating of a fresh sample of MnS
A fresh sample of MnS which did not give a green colour when treated with the dihydrochloride reagent was heated with a bunsen burner for 10 minutes. On treat•
ing this sample with the CeHBN.2HCl in CHC13 reagent, the solution turned light green. The sample after heat• ing for a longer time turned the reagent solution a darker green.
iv. Effect of oxidizing agents
The CBH5N.2HC1 in CHC13 reagent was allowed to -28- react with manganese metal. Samples of the resulting
solution were treated with CeH5N.HN03, HN03 (cone),
KsCrs07, NaOCl, and HS0S. The addition of any of these oxidizing agents caused the formation of the familiar green colour. They also caused the colour to appear in
solutions of the reagent and MnCls or MnS04. Heating the solution of manganese metal in the reagent did not have any apparent effect.
v. CBHBN.2HC1 in CHC13 on Mn(0Ac)3
The Mn(0Ac)3 crystals prepared as described in
(4.a. ) reacted with the reagent to give a dark green colour.
vi. Isolation of the substance conferring
the green colour to the CBHBN.2HC1 in CHC13
CBHBN.2HC1 in CHCl^ solution was reacted with an
excess of Mn02. The dark green solution which resulted was poured into an excess of alcoholic HCl solution3'1 and on the addition of dry ether a dark green crystalline precipitate resulted which was filtered and dried in a vacuum desiccator. The melting point was 100-101°C.
It was found that the solutions used for isolating the crystals had to contain an excess of HCl or light green to whitish crystals resulted. The crystals appeared to
* The alcoholic HCl solution was prepared by bub• bling dried HCl gas from a Kipp generator through abs• olute alcohol. -29- be quite hygroscopic and dissolved in moist air with a loss of colour. It was found that the original sol•
utions of the green substance in CBHSN.2HC1 in CHC13 also lost their colour on standing. The green cryst• als were very soluble in water, and on first dissolving a pink colour appeared momentarily.
vii. Analysis of the green crystals
The soluble chlorine content was analyzed by
Fajan s Method and the manganese content was determ-
33
ined gravimetrically as MnNH4P04.Hs0 . The remainder was assumed to be pyridine.
% * Atomic Weight Ratio
Mn 11.5 55 = .21 1
CI 37.6 35.5 = 1.06 5
C5H5K 50.9 79 = .64 3
vili. Limit of visible colour
Solid KMn04 (.0.2 gm. ) was dissolved in 10 cc. of pyridine. This solution was added dropwlse to 2 cc. of
the reagent of CBHBN.2HC1 in CHC18. The green colour was just visible after the addition of 0.02 cc. of the 5
KMn04 solution or 5 X 10" gm. of manganese,
ix. Colourometric analysis
A Coleman Spectrophotometer was used for the
colourometric work. As a reference standard, 10 cc.
of a saturated solution of pyridine dihydrochloride in -30
chloroform was used0 The spectrophotometer was adjusted
so that for this,reference solution concentration C = 0,
and transmittance T = 0. KMn04 dissolved in pyridine was added to a similar solution of reagent until a medium green colour was obtained. A Specttfal-Transmittance
curve was made (Wavelength in millimicrons vs % Trans• mittance) to determine the portion where T was essent•
ially constant (Figure III). This region, at 665 milli• microns was used in all subseuent work.
A solution of KMn04 in pyridine was added drop- wise to a saturated chloroform solution of pyridine
dihydrochloride and the percent transmlttanee was record•
ed as was, the volume of solution containing KMn04. The
concentration of KMn04 was determined, accurately by
titration against a standard solution of Ass03. The
volume of an average drop was measured, and hence the
concentration of manganese in grams per dribp was known.
All the manganese was assumed to exist as the green
complex and KMn04 solution was added till the percent
transmittance dropped to a small value (Figure IV).
Another run of the percent transmittance and conc•
entration in grams of manganese was obtained by adding
a known amount of manganese metal dissolved in a known
volume of saturated solution of C5HBN.2HC1 in chloroform
to 10 cc. of a saturated chloroform solution of
CBH6N.2HC1 in which excess C5HBN.HN03 was dissolved.
-33- This reagent produced the green colour on the addition of the manganese metal (Figure V). x. Variation of colour with time
KMn04 in pyridine was added to the pyridine dihydrochloride reagent in chloroform to conform to 6 55 X 10"* gm. of manganese in 10 cc. of reference sol• ution and the variation of percent transmittance was studied with respect to time (Figure VI).
c. Reaction of CBHBN.2HC1 in CHCla with HgS i. Isolation of mercury salt
A solution of CBHBN.2HC1 in CHC13 was allowed to react with excess HgS. A vigorous reaction ensued,
and the resulting solution was boiled to remove H3S, which if present, was found to reprecipitate the HgS on attempting to Isolate the mercuric salt. The boiled solution was poured into excess ethanol from which a white crystalline precipitate was obtained, which, after filtration and after drying in a vacuum desiccator had a melting point of 123°C. ii. Analysis of the mercury salt The crystalline mercuric compound was soluble in water. The chlorine content was determined by Fajan1s 88 Method . Mercury, if left in solution, led to erroneous results due to the formation of the soluble but non-
ionized HgCl8. It was removed by prepitation with
NasC03 and filtration of the solution. The amount of
FIGURE VI. VARIATION OF COLOUR OF GREEN TRIVALENT MANGANESE COMPLEX WITH TIME. -36-. mercury was determined gravimetrically as
LCuCNHsCHaCHgNHg )sHHgI4i and the remainder was assumed to be pyridine.
io * Atomic Weight Ratio
Hg 39.6 201 = 0197 1 CI 28.0 35.5 = .788 4
CgHgN.H 32e4 80 = e40§ 2
iii. A compound was made by dissolving HgCl8
in CeHBNoHCl. Aprecipitate was also obtained fuom ethanol which had a melting point of 123°C. A mixed melting point with the first mercuric salt also was 123°C. -37
III •* DISCUSSION
Pyridine salts as rectlon media.
a. Type of reaction most desirable
It was decided that the most suitable reactions to investigate would be those of metals and sulphides.
Sulphides were chosen because acid reactions with them
would produce H2S which is given off as a gas and would thus serve, not only to aid the reaction in going to completion, but also to act as an easily observed criterion for a reaction occurring. Furthermore, it was decide to avoid those reactions in which water might be formed as the presence of water might make it dif- 2 ficult to study the reaction products .
b. Examination of the pyridine salts prepared
Pyridinium trichloroacetate could not be used as it decomposed on melting. Pyridinium oxalate and pyr• idinium nitrate often produce water as one of their reaction products and were therefore discarded. Many of the other salts used as acids, such as pyridinium thiocyanate were too weak to dissolve many of the metals and their salts. It was not possible to prepare pyr• idinium fluoride with the equipment available. Pyridin• ium fluoride might be a good medium, but would have the disadvatage of being hard to handle. It should be pos- -38- sible to prepare pyridinium fluoride by reacting pyr• idine with liquid HF in a stainless steel container.in the ratio of one mole of pyridine to about six of HF, fol• lowed by evaporation of any excess of the low boiling
HF, This excess of HF is necessary as amines form com• pounds of the type B.4HF, where B is a primary, second- 6 ary, or tertiary amine, Berliner and Hann propose the structure:
ii: B H-.F-H-.F: ii
It was found that the pyridine salts most suit• able for investigation as solvent systems for metals and sulphides were pyridinium chloride and pyridine dihydrochloride. These two salts reacted with many metals and sulphides, were reasonably easy to obtain in a pure state, and had the advantages of pyridine
salts in general, that is, the stability of the pyridine nucleus and the existence of many readily isolated crystalline derivatives. Furthermore, there was no possibility of the formation of water as a side prod• uct and the reaction should be relatively simple.
Methods of using CBHKM.HC1 and CKHKN.2HC1.
a. Molten state
Reactions with molten pyridinium chloride have been
carried out with a number of metals, metal oxides, and -39- S 9 8 6 ' 2 8 metal chlorides . Ho*ever, although pyridinium
chloride melts at the relatively low temperature of l44°C., reactions with metals may result in side products
as were evidenced with Al, Mg, and Zn. With the except•
ion of these three metals it was an excellent solvent medium for many metals and sulphides,,
Pyridine dihydrochloride was also a fairly good medium for the substances used. However, both pyridine
salts had certain disadvantages. They were both qtiite
hygroscopic and, since the presence of water was to be
avoided, this was a decided disadvantage. Furthermore,
although the temperatures at which the two hydrochlor•
ide salts melted were low, it would be more convenient
to find some system where heat was not necessary to
produce a medium which would react with the metals and
sulphides used.
b. Room temperature systems
The disadvantages mentioned in the previous section
acould be overcome by the use of chloroform solutions of
pyridinfeum chloride and pyridine dihydrochloride. Both
salts were relatively soluble in chloroform and both
still dissolved many metals and sulphides. The second•
ary reactions of Mg, Al, and Zn noted with molten pyr•
idinium chloride were not evinced with either solvent.
The reactions with the dihydrochloride solution were
more vigorous than the chloroform solution of the mono- -40- hydrochloride and appeared to be only slightly less reactive than molten pyridinium chloride. The presence of the additional mole of HCl thus influenced the react• ivity of the solvent and resulted In reactions with some salts of manganese that did not occur with any of the other reaction mixtures. The reactions in general may be expressed in two steps:
2C6H6N.HC1 + MeS - 2CeH6N + MeCl8 + H8S and, if the formation of a double salt is possible
2C5H5N.HC1 + MeCl8 -—* (CeHBN.H)8MeCl4 where Me Is an^ example of any divalent element whose sulphide reacts.
It might be expected that only those sulphides reacting with aqueous hydrochloric acid would dissolve in the systems studied. However, the stability of the pyridine complexes must be considered as well as the difference in ionizing and dissociating powers of the nonaqueous systems. Mercuric chloride was found to dis• solve quite readily in the chloroform solution of pyr• idinium chloride. In the water - hydrogen chloride system the sulphide ion concentration cannot be reduced below that which will allow the mercuric salts to stay in solution. However, in the chloroform solution, the stability of the mercury complex, and the inhibition
of the ionization of the H8S, allow HgS to be dissolved
by CBHBN.HC16 The nature of CBHeN.2HCl In CHC13.
There is very little reported in the literature
concerning the nature of the compound CBHBN.3HC1. is;
Kaufler and Kunz , who first prepared the substance, stated that HCl polymerized in the same manner as the
hydrofluorides (H3F8, H3F3). In their subseuent work, they found that the stability of the dihydrochloride was essentially limited by the degree of alkylation and.that only tertiary and quaternery bases formeithe dihydro- chlorides. No trihydrochlorides Mgre reported. The dihydrochlorides were simply postulated as:
x
C y NH][C18H] z Since this compound was prepared, there has been no further suggestions as to why tertiary bases such as pyridine should take on a second molecule of HCl.
In order to determine if there existed any type
of bonding of the second HCl molecule to the CBHeN.HCl molecule in chloroform solution, the wavelength region
X = 3300-5300 A was investigated on a Beckmann Spectro• photometer, The saturated chloroform solution of pyr• idine dihydrochloride on dilution showed that the plat• eau A in Figure I was the same peak (Figure II) as was given by the monohydrochlorlde solution. On the basis of the similarity of the plots, such possible structures as pi bond formation are ruled out, and in the chloroform -4 ab•
solution, the CBHBN..2HC1 can be considered as CBHBN.HC1
+ HCI (the presence of HCI does not effect the absorb- ency). This seemed quite reasonable as the evolution of HCI fumes from such chloroform solutions was very noticeable.
Side reactions.
a. Possibility of formation of blovridvls
Pyridine has a structure analogous to that of benzene. The introduction of the nitrogen atom into, the ring in\place of a carbon atom results in a compound with its own particular reactions, yet maintaining a certain structural relationship to the "parent" benzene. The resonance structures contributing most to the stability of the pyridine ring can be represented as:
Hydrogenation, if it did occur, would be very difficult due to the inertness of the pyridinium ion. This is explained by assuming that the natural electron attract• ion of the nitrogen atom in the pyridine ring is enorm• ously enhanced in an acid solution where the pyridine exists as the positive charged pyridinium ion VI. The effect of the positive ion could be represented accord• ing to the designation of the English school as VII, -43* which indicates that the electron attraction of the positively charged nirogen atom would reduce the electron density in the 2 and 4 positions and therefore would enhance the polarization Indicated by the III, IV, and
V resonance states.
M TEL
Since hydrogenation could be represented as attack
+ by H3 , it is obvious that hydrogenation of an already : positive charged pyridinium ion would be difficult and that, when it did occur, it would attack, not the 2 and
4 positions of low electron density, but would attack the position which is relatively unaffected by the quat• ernary nitrogen atom, the 3 position.
In practice, no bipyridyls were found and the side reactions must be accredited to some other type of reaction.
b. Ring Cleavage
Ring cleavage reactions with pyridine are possible as the pyridine ring offers a nitrogen atom with an un• shared pair of electrons as a point of attack. This [email protected] not the case with pyridinium compounds. However, with certain special derivatives the normal pyridinium -44- resistance is lost. It is suspected that the side react• ions which occur with pyridinium chloride and Zn, Mg, and Al may be of this type.
One sueh special compound is 4-pyridyl-pyridinium dichloride:
Ring opening in this compound should follow the proposed reaction for 2,4-dlnitrophenylpyridinium chloride in 40)41)43 alkaline solution . 4-pyridyl-pyridinium di• chloride undergoes a similar reaction with the production of a red-brown product quite similar to that obtained in the reaction of pyridinium chloride plus Zn. The main evidence in favour of an open chain structure with the
4-pyridyl-pyridinium dichloride and 2,4-dlnitrophenyl• pyridinium chloride Is their colour. Since the orig• inal structures were colourless, the coloured compounds suggested a form with an increase in the conjugated sys• tems to account for an increased absorption of visible light. If ring cleavage occurs in the pyridinium chloride plus Zn reaction, it does so in the fefsence of any oxy• genated substance,, This eliminates the enol type struct• ure as is found with the two examples of ring opening discussed. Because of the coloured reaction product there should exist some increased conjugation.- Further•
more, the decolouration of a dilute KMn04 solution sug• gested an open chain with at least one double bond. This seemed substantiated by the loss of colour of a solut•
ion of the unknown product when treated with 30$ H30s and allowed to stand. If the side ohain postulated was oxidized, this would account for the loss of colour.
It had been hoped that a study of the reaction products of quinoline hydrochloride and metals would offer some clue as to the nature of the reaction. Further• more, the large increase in yield of the quinoline hydro• chloride side product was an additional advantage. At this point however, research was stopped.
c. Limitations due to side reactions
The only side reactions observed were those which occurred at high temperatures w\ifchpyridinium chloride and only with the electropositive metals Zn, Mg, and Al.
These reactions produced only a very small quantity of reaction products and were assumed to occur to only a slight extent. In the other media no side reactions were observed with either metals or salts. Since ^matty metals -46-
and sulphides are soluble in some of the solvent sys•
tems studied, they offer excellent media for solution.
The solvent system of special note because of its high reactivity and use at room temperature was pyridine
dihydrochloride in chloroform,,
Reactions with possible analytical-application,
a. Trivalent manganese
i. Formulation of the compound
(05HsN.H)aEMnClBiCeH6N
It was noticed that C5H5N.2HC1 in CHC13 gave dis-.
similar reactions with different samples of MnS. A
freshly prepared sample dissolved to give a colourless
solution, whereas an old sample resulted in a dark green
colour. Freshly prepared MnS, when heated on exposure
to air, and again treated with the chloroform solution
also gave the green colour. Manganous salts such as
MnCls and MnS04 did not give the great colour in the
chloroform reagent, but it appeared on treating the res'* ultant solutions of these salts with an oxidizing agent.
All the foregoing evidence seemed to indicate that the
green colour was due to the existence of an oxidation
state e# greater than +2,
MnOa reacted very vigorously with the reagent and
the evolution of chlorine was noticeable, suggesting
that the oxidation state of the resulting manganese
compound was less than +4. The presence of manganese in the trivalent state seemed <#ite evident. Confirm• ation was obtained when a colourless solution resulted
on reacting Mn(0Ac)3 with the pyridine dihydrochloride
in chloroform whereas Mn(0Ac)3, resulted in a green sol• ution on identical treatment. Analysis of the isolated compound suggested the formulation of the complex
(C6HeN.H)8CMnCl63c6H6N.
Calculated Theoretical
CI 37.6$ 37.6$
Mn 11.5$ 11.6$
This particular structure was suggested as the complexes of trivalent'manganese are confined to the unusual type
MaCMniBl, often with a molecule of water which presumably
completes the coordination number of 6; the type M3[MnX63 87
is not known to occur . The compound isolated in this
case had a molecule of pyridine to complete the coordin•
ation number of 6 instead of a molecule of water.
ii. Explanation of reactions of manganese
salts
1. It has been shown that when MnS was heated oxidation takes place In such a manner that mang- anese^in an oxidation state of greater than +2 as the
heated MnS reacted with the CBHBN.2HC1 in CHC13 to give
the green coloured trivalent manganese complex. -48-
2. Pure crystalline MnC03, is pink. -It
slowly darkens in air through oxidation. It evolves C0S leaving MnO which is readily oxidized in air to higher S 0
oxidation states, e.g, MnsOa and Mns04 , Since the
sample used was not a fresh one, the presence of Mn303
and Mn304 account for the green colouration.
3. The colour with manganous nitrate
is to be expected due to the presence of the nitrate
ion. It has been shown that oxidizing agents such as
HN03 cause the formation of the green complex from stable
divalent manganese salts.
4. MnOl3, MnS04, Mn(0Ac)8 did not give
a green colour as an oxidation state of manganese greater
than +2 was not formed.
iii. Determination of trivalent manganese.
by colourometric analysis of the green complex
The Lambert-Beer Law (1^ = IQIO"*^01) is the only -
rational means for translating photometer readings to
expressions of the corresponding concentration of the
sample. A simple form for this expression is:
. C = -K log T where C is the concentration, K is a proportionality
constant and T is the transmittance (T = It/lo)-» ^
the Lambert-BeBr Law is obeyed, the Concentration-
Transmittance graph of this relationship plotted on semi•
log paper, is a straight line intersecting the point -49-
0=0, T = 100$. It was necessary to show that solutions of the green trivalent manganese complex obeyed this law in order to determine the T of one solution of known concentration and then draw a straight line intersect• ing this point and the point C = 0, T = 100$. Such a straight line then represents a valid Concentration-
Transmlttance graph and will allow the contration of any other solution to be determined by measuring only the percent transmittance.
The Lambert-Beer Law required that the T measure• ment be made with monochromatic light and at a wave• length corresponding to a region of the constituent's
Spectral-Transmlttance curve where T was essentially constant. The reference selected must be such that
C = 0 when T = 100$ and the nature of the sample sol• ution must be such that its T responds only to changes in C. With these limitations the flat portion best chosen was 665 millimicrons (Figure III).
Figure IV, obtained by the reduction of KMnO*, showed that the Lambert-Beer Law was obeyed in the region of 8 X 10* gm. of manganese per 10 cc. of solution down _ e to 3 X 10 gm. of trivalent manganese in the same volume of solution. Figure V is a similar plot, but obtained by forming the trivalent complex by oxidation of a manganous solution. Concentrations of trivalent -4 manganese greater than 8 X 10 gm. in 10 cc. of solution -50-
can be determined by suitable dilution. It must be noted, however, that the colour represents not only any manganese originally present in the trivalent state, but any manganese of higher valence which may be reduced to the trivalent state, and divalent manganese which may become oxidized.
Figure VI represents the variation of colour of a 6
55 X 10" gm. trivalent manganese complex per 10 cc. of reagent solution with time. The plot shows that the time may be a very important factor and that any colour•
ometric analysis based on the use of the green trivalent manganese complex will have to take into consideration
the fading of the colour with time.
b. The mercurv-pyrldine complex
i. Formulation of the compound (CBH5N.H)aHgCl4
Analysis of the isolated mercuric compound sug•
gested that the formula of the compound was (CBHBN.H)aHgCl4.
Calculated Theoretical
Hg 39.6$ 39.9$
Cl 28.0$ 28.2$
The mixed melting point of this compound prepared from
pyridinium chloride in chloroform with HgS and the
compound isolated in a similar manner from HgCla plus
C5HBN.HC1 was 123°C. This value agreed with that of
13 Grossmann and Hunseler for the above compound. -51-
iio Application of solubility of HgS
The one application of the solubility of HgS in
the reagent solution CBHBN.HC1 in CHC13 that suggested
itself immediately was in radiochemical separations.
HgS is an excellent earrler for many radioactive isotopes,
and the ready solubility of HgS in the reagent solution
used here, may simplify separation techniques.
6„ Conclusions.
It is evident that reactions in media such as
molten pyridinium chloride or pyridine dihdrochloride
in chloroform are different in many respects to similar
reactions carried out in water. The strong bonding of
the hydrogen atom to the nitrogen atom of the pyridine
ring allows many salts to be heated to high temperatures
without decomposing.
The hydrochlorides of pyridine have been shown -
to be strong enough as acids to act as good dissolving
reagents. They have the further advantage of forming
complexes which are stable. Also many of these complexes
can be isolated as crystalline derivatives at room temp•
erature. Thus, these hydrochlorides of pyridine offer
media in which the sometimes disadvantageous ionization
of water may be circumvented. The formation—of Many
complexes which would dissociate in water can occur and
oaa be isolated. -52-
IV - BIBLIOGRAPHY
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