SOLUTIONS IN DIFLUOROPHOSPHORIC ACID
' -by William Reed
A THESIS SUBMITTED IN PARTIAL FULFILMENT OF
-THE REQUIREMENTS FOR THE DEGREE OF
DOCTOR OF PHILOSOPHY
in the Department
of
Chemistry
We accept this thesis as conforming to the required standard
THE UNIVERSITY OF BRITISH COLUMBIA
January 196S
0 William Reed 1968 In presenting this thesis in partial fulfilment of the requirements for an
advanced degree at the University of British Columbia, I agree that the
Library shall make it freely available for reference and study. I further
agree that permission for extensive copying of this thesis for scholarly
purposes may be granted by the Head of my Department or by his represen•
tatives. It is understood that copying or publication of this thesis for
financial gain shall not be allowed without my written permission.
Department of
The University of British Columbia Vancouver 8, Canada
Date February 27, 1968 (ii)
ABSTRACT -
The physical and inorganic chemistry of solutions in difluorophosphoric acid, HPO^Fg, has been studied, as part of a general study of solutions in non-aqueous protonic solvents.
Difluorophosphoric acid is a colourless, associated liquid which might be expected to have solvent properties similar to those of other protonic systems such as H^O, I-^SO^ and HSO^F.
However, electrical conductivity studies of solutions of various electrolytes and nuclear magnetic resonance studies of solutions of alkali metal difluorophosphates indicate that the acid is a poor solvent for electrolytes and that ion-pairing is probably extensive.
Acid-base behaviour in HPO2F2 has been extensively in• vestigated. Compounds which behave as bases in this system in• clude metal difluorophosphates, chlorides, nitrates and carbonates, organic amines, and some organic nitro-compounds and carboxylic acids. Inorganic molecules such as F^SO^, HSO^F and SbF^ behave as acids. Reaction between an acid and a base in HPO2F2 commonly result in the formation of an insoluble salt. The reaction between KPO2F2 and SbF^, for example, has been used to prepare the new compound KSbF^P02F2«
To further investigate the factors affecting acid strengths, cryoscopic and electrical conductivity studies of various inorganic oxy-acids were carried out in nitrobenzene, as solvent. The acids H^SO^, HSO3F and HPO2F2 appeared'to be vir• tual non-electrolytes in nitrobenzene, vrith ^SO^ apparently exhibiting some polymerization. (iii)
TABLE OF CONTENTS •
PAGE
CHAPTER I General Introduction ' 1
1.1 Properties of difluorophosphoric acid 1
1.2 Acid-base behaviour in protonic solvents 4
1.3 Outline of present work 6
CHAPTER II Solutions of Metal Difluorophosphates 8
2.1 Introduction 8
2.2 Experimental 8
A) Preparation and purification of materials 8
i difluorophosphoric acid ii metal difluorophosphates
B) Electrical conductivity 10.
C) Nuclear magnetic resonance 17
D) Viscosity 18
E) Density 19
2.3 Results and discussion 19
A) Electrical conductivity 19
B) Nuclear magnetic resonance 30
C) Density 41
D) Viscosity 1+6
CHAPTER III Miscellaneous Bases 48
- A) Organic solutes 48
3.1 Introduction 48
3.2 Experimental 48
A) Electrical conductivity 48
B) Preparation and purification of materials 48
3.3 Results and discussion 49 (iv) • . PAGE
B) Inorganic solutes . 53
3.4 Introduction 53
3.5 Experimental 54
3-6 Results and discussion " 54
CHAPTER IV Acids and Acid-Base Reactions 62
4.1 Introduction 62
4.2 Experimental 63
A) Preparation and purification of materials 63
B) Electrical conductivity 64
C) Nuclear magnetic resonance 65
4.3 Protonic acids: results and discussion 65
4.4 SbF^ solutions: results and discussion 72
4.5 Studies on KSbF5P02F2 95
CHAPTER V Nitrobenzene. Solutions 102
5.1 Introduction 102
5.2 Experimental 103
A) Cryoscopy 103
B) Electrical conductivity 108
C) Preparation of materials 109
5.3 Results and discussion 109
A) Fluorosulphuric acid solutions 110
B) Sulphuric acid solutions 117
C) Pifluorophosphoric acid solutions 121
5.4 Conclusion 121 CHAPTER VI Summary and Suggestions for Further Work 123
6.1 -Summary 123 6.2 Suggestions for further work 125
BIBLIOGRAPHY .127 (v)
LIST OF TABLES . •
TABLE PAGE
1. Physical Properties of Difluorophosphoric Acid t 2
2. Specific Conductivities of the Alkali and Some Alkaline 20 Earth Metal Difluorophosphates at 25°
3. Equivalent Conductivities of Some Difluorophosphates at 25° . 23
4. Specific Conductivities of Some Potassium Salts in Various Solvents 26
1 19 1 5. H, F and ^ P Chemical Shifts for Solutions of MP0„F9
in HP02F2 32
6. Densities and Viscosities of Some Solutes in HP02F2 at 25° 42
7. Specific Conductivities of Some Organic Bases in KPOoFo at 25° . 50 • #. Specific Conductivities of Various Electrolytes in HPO F at 25° 55 2 2
9. Specific Conductivities of Some Acids in HP02F2 at 25° 66
10. 19F.and 31p Ch emical Shifts and Coupling Constants for Some Complex Antimony-Fluorine Species 82
11. Infrared Spectra .of Various Inorganic Fluorine Compounds 96
12. Infrared Snectrum of Gaseous Products from the Decomposition
of KSbF5P02F2 100
13. Cryoscopic Measurements in Nitrobenzene 111
14. Specific Conductivities of Some Electrolytes in Nitrobenzene at 25° 114 15. Equivalent Conductivities of Some Electrolytes in Nitrobenzene at 25° 119 (vi)
LIST OF FIGURES
FIGURE PAGE
1. Difluorophosphoric Acid Distillation Apparatus 9
2. Electrical Conductivity Cell - 11
3. Injector, used for Solute Additions to the Conductivity . Cell ' 13
4. Microburette, used for Solute Additions to the Conductivity Cell 16
5. Specific Conductivities of Some Difluorophosphates at 25° 21
6. Equivalent Conductivities of Some Dif luorophosphates at 25° plotted against the Square Root of the Ionic Strength 25
7. N.M.R. Chemical Shifts for MP02F2 in HP02F2 34
19 8. F N.M.R. Chemical Shifts for MPOgFg in HP02F2 35 31 9. P N.M.R. Chemical Shifts for MP02F2 in HP02F2 36
10. Densities of Some Metal Difluorophosphates in HPO^Fo 44 at 25°
11. Specific Conductivities o-f Some Organic Bases in HP09F9 at 25° 51
12. Specific Conductivities, of Various Electrolytes in HP02F2 at 25° 56
13- Specific Conductivities of Various Potassium Salts in
HP02F2 at 25° 59
14. Specific Conductivities of Some Acids in HP02F2 at 25° 08
15. Acid-Base Titrations in HP02F2 at 25° 70
16. - Acid-Base Titrations for SbF against KPO F in HPO F at 25° 5 . 1 Z 1 75 19 17. F N.M.R. Spectrum of a 3.63 molal SbF£./HP0oFo Solution 5 at 30° ^ d. 7g
19 18. F N.M.R. Spectrum of a 3.63 molal 3bF_/HP0oF Solution at -65° 5 2 2 79
19. 19F N.M.R. Spectrum of a 7.0 molal SbF /HPO F Solution
at 30° 5 2 2 80 (vii)
FIGURE . PAGE
20. 1^F N.M.R. Spectra of the P-F Region for a 2.31 molal
SbF5/HP02F2 Solution at 30° and -70° 84
21. 19F N.M.R. Spectrum in the Sb-F Region for a 2.31 molal '
SbF5/HP02F2 Solution at -70 (peak M) 85
22a) X9F N.M.R. Spectra in the Sb-F Region for a 2.31 molal •
SbF5/HP02F2 Solution at -70° (peaks K, L and P) 86 22b) ^F.N.M.R. Spectrum in the Sb-F Region for a 2.31 molal
SbF5/HP02F2 Solution at -70° (peaks N & 0) 87
23. X9F High-Resolution N.M.R. Spectrum of the P-F Region for
a 2.31 molal SbF5/HP02F2 Solution at 30° 92
24. 51P N.M.R. Soectrum of a 2.62 molal•SbFr/HPO-F- Solution at 30° ? 93 25. Details of the31 P N.M.R. Spectrum of a 2.62 molal SbF./ 5 HP02F2 Solution at 30° 94
26. Vacuum Line used for the Decomposition of KSbF^P02F2 98
27. Cryostat, used for Nitrobenzene Solutions 104
28. Depression of Freezing Point (AT) for Various Solutes in Nitrobenzene 112' 29. Specific Conductivities of Some Solutes in.Nitrobenzene at 25 116
30. Equivalent Conductivities of Some Electrolytes in Nitrobenzene at 25° 118 (viii)
ACKNOWLEDGMENTS
The author wishes to express gratitude to Dr. R. C.
Thompson who first suggested the problem, and under whose guidance the work was done.
Thanks are due to Mr. S. Rak who constructed the glass apparatus, to Mr. R. Burton who operated the H.A..100 n.m.r. spectrometer, to Mr. R. Wolfe who assisted with the operation of the platinum resistance thermometer and finally to Mr. L.
Neering for many helpful suggestions.
The generous gift of difluorophosphoric acid by the
Czark-Mahoning Chemical Company is also gratefully acknowledged. CHAPTER I
'• ..' -General Introduction
At the turn of the century solution chemistry was largely,
concerned with reactions carried out in aqueous media. During the past five decades, however, studies on a variety of non• aqueous solvents have resulted in the development of many new
solvent systems. The experience so obtained has greatly broad•
ened the scope of synthetic chemistry and has considerably in•
creased the understanding of the physical and chemical properties
of solutions. Although the number of solvents which have been
investigated is very large, extensive and systematic studies on the physical properties of solutions have been limited to rather few solvents, notably HF, H^SO^ and NH^ and to a smaller extent
-HSO^F, S02 and HC1.
The purpose of the work described in this thesis was to
study the properties of solutions in anhydrous difluorophosphoric
acid and to investigate the possibility of HP02F2 as a prepara•
tive medium. At present, the range of suitable and readily
available fluorinated solvents of use for the preparation of
fluorides is limited essentially to hydrogen fluoride, fluoro•
sulphuric acid and bromine trifluoride and it is hoped that
HP02F2 will extend this range.
1.1 Properties of difluorophosphoric acid
The literature regarding difluorophosphoric acid has been
2 reviewed recently >3,4 and as little fresh information has ap•
peared since these reviews were written, the usual historical 2
review of the solvent will be dispensed with. However-, a sum•
mary of the physical and chemical properties of HPO2F2 relevant
to this work will be attempted. Difluorophosphoric acid is a
mobile, colourless, clear liquid which fumes strongly in moist 2 air ; some of its physical properties are listed in Table 1.
Table 1 .
Physical Properties of Difluorophosphoric Acid
Property Value (reference 3) Value(this work)
Melting point °C: (i)-91.3-1.0 (ii)-96.5-1.0
Boiling point °C: (i) 108-111 (ii)115.9
42-43' at 75mm. 50-52 at 100mm.
Liquid density {&/ml): d^5 = 1.583 dj?5 = 1.583 viscosity YJ (poise): 0.0519 at 26.01° 0.055-.003 at 25.0° heat of vaporization (cal./mole) : (i)7925, UD9125 (iii)9360 Trouton constant: (i)20.4, (ii) 23.7 (iii).24.6
It is a monobasic acid in water and is slowly hydrolysed with the 5 formation of monofluorophosphoric acid ,
HP02F2 + H20 —* HgPO^F + HF (1.1).
Trotter et al. carried out the first detailed study of the P02F2~ ion in their structure analysis of potassium difluorophosphate.^
It was found that there is distortion of the valency angles in the difluorophosphates from the regular tetrahedral value. The
molecular shape of monomeric HP02F2 may, therefore, be assumed to be also approximately tetrahedral. The degree of association of molecules in a liquid determines to no small extent its solvent properties and the evidence that HPO2F2 is a strongly associated liquid is extensive.
In the series of compounds shown below the molecular weight of each molecule in each column Is approximately the same i.e. a fluorine atom has been replaced by a hydroxyl group. -
S-compounds boiling point (1 atm) P- compounds boiling point (1 atm);
S02F2 -55.4 POF^ -39.8
H0S02F 162.7 HOPOF, 116 (with decom• position)
(HO)2S02 317 (with decom• (H0)2P0F (cannot be distil position) led , some decom• position at 180°)
(H0)3P0 213 (-|H20)
The large change in boiling point is attributed to molecular 7 association due to hydrogen bonding . In the sulphur series of compounds the very high boiling point of H-SO compared to that 8 d k of HSOoF is consistent with a higher degree of molecular associa• te tion in the former liquid. It would appear on the basis of its boiling point, therefore, that difluorophosphoric acid is exten- 9 sively associated. Lenskii et al. report a value for the vis• cosity of HPO F of 5.31 centipoise (c.f. I.56 cp. for HSO F),10 22 11 3 which is further indication of solvent association. This associa 1 p tion was further studied by Stafford et al. who examined the infrared spectra of a number of hydroxylic acids in the liquid and vapour states and found the OH stretching vibration shifted to lower frequencies on going from the vapour to the liquid. 4
It was concluded that in the series of compounds HNO^, HCIO^,
HSO3F, CH3SO3H and HP02F2 the strength of hydrogen bonding is in•
dicated by the extent of the vapour-liquid shift and that as the
OH frequency shift is greatest for HP02F2 it probably forms the
strongest hydrogen bonds.
The Trouton constants of many non-associated liquids are
of the order 21.5,^ while associated liquids commonly exhibit
somewhat greater values. Consistent with this, the Trouton con•
stant of HP02F2 is approximately 24. Finally, evidence for hydrogen bonding may be obtained from the n.m.r. chemical shift of protonic liquids.1^ In the present work it was found that
the resonance of HP02F2 occurs at -3.35T lower field even than that found for HgSO^ (-1.6T).
1.2. Acid-base behaviour in protonic solvents
It is convenient to discuss the chemistry of solutions in protonic solvents in terms of acid-base behaviour and a few com• ments concerning definitions of acids and bases would be. relevant at this point. It is known that water is not unique in its ability to act as an ionizing solvent and as a medium for acid-base behaviour. Analogies among protonic solvents become apparent when 15 the auto-protolysis reaction of each is considered. Although such ionization is ordinarily comparatively small, conductivity measurements have shown its existence in a number of instances. In the table below the autoprotolysis reactions of various sol• vents are listed. 5
Solvent
+ H20 H20 + H20 ^ H^0 + 0H-
+ NH3 NH^ + NH^ ^ NH^ + NH2~
+ HF HF +• HF ^ H2F + F~
+ H2S04 H2S04 +' H2S04 ^ H3S04 + HS04"
HSO^F HSO^F + HSO^F ^ HgSC^F* + SO^F"
Based upon such considerations several definitions for acids and bases have been advanced in terms of the parent solvent. 16,17,16*
They can be combined to admit as an acid any material giving, either by direct dissociation or by interaction with the solvent, the cation characteristic of the solvent and as a base any material giving the anion characteristic of the solvent. Neutralization, in terms of the general theory of solvent systems, amounts to the combination of the solvent cation with the solvent anion to pro• duce the solvent.
By analogy with other protonic solvents acid behaviour in
HP02F2 may be defined as the reaction of a solute to produce H- the acidium ion, H^PO^F^ and base behaviour as the reaction to
produce the P02F2~ ion. These ions are the characteristic ions produced by solvent autoprotolysis.
+ 2HP02F2 ^ H2P02F2 + P02F2~ (1.2)
There is some evidence (to be given later) for this reaction in
HP02F2, however, the definition of acids and bases in this solvent is actually independent of whether or not this reaction takes place to a measurable extent. Difluorophosphates which ionize
to give P02F2~,
+ MP02F2 ^ M + P02F2" (1.3) and molecules which are protonated,
+ B + HP02F2 BH + P02F2" (1.4) are bases in this system. Correspondingly, acids are molecules
which protonate HP02F2,
+ HA + HP02F2 — H2P02F2 + A" (1.5) or molecules which accept difluorophosphate anions,
+ A + 2HP02F2 ^ AP02F2" + H2P02F2 (1.6)
1.3. Outline of present work.
Chapter II of this thesis is concerned largely with the detailed investigation of the nature of solutions of the alkali and alkaline earth metal difluorophosphates in difluorophosphoric acid. This study has closely followed that reported by Gillespie 19 and co-workers 7 on solutions of metal hydrogen sulphates in sulphuric acid. Using a variety of physical measurements these authors were able to elucidate the nature of these solutions.
From electrical conductivity, transport number and freezing- point depression measurements they were able to show, for example that the metal hydrogen sulphates are fully dissociated electro• lytes, the degree of solvation of the alkali metal cations de• creases with increasing cation size and the hydrogen sulphate anion conducts electricity by a proton-transfer mechanism.
Protonation reactions of organic molecules in solvents of
20 10 21 high acidity such as HF, HSO^F and H2S0^ have been studied extensively. Chapter III describes electrical conductivity studies on protonation reactions of organic bases in HP02F2.
Electrical conductivity studies on solutions of miscellaneous 7 inorganic compounds are also described in this chapter. 3
Although HPC^Fg is a strong acid in H20 , studies of solu- 77
H F t0 be a weaker acid tions of HP02F2 in H2S0^ have shown P02 2 than both HgSO^. and HSO^F and therefore a wider range of solutes
which would behave as acids in HP02F2 should be available.
Chapter IV is concerned with acid behaviour in HP02F2 and acid- base reactions in this solvent.
Factors which determine the relative orders of acid 22 strengths are not well understood. Previous work has estab•
lished the order HSO^F > H2S0^ > HP02F2; however it is signifi• cant that this order is established only for the case where the bulk solvent is a protonic medium. In an attempt to determine the relative proton-donating abilities of these acids in non- protonic media, solutions of the acids were studied in nitro• benzene as solvent and this work is described in chapter V. CHAPTER II
Solutions of Metal Difluorophosphates
2.1 Introduction
The alkali and alkaline earth metal difluorophosphates may be expected to act as strong bases according to (1.3) when
dissolved in difluorophosphoric acid. It is not possible, however, to predict the degree of dissociation of these salts into free ions, particularly as the dielectric constant of HPC^Fg is not
known. In an attempt to obtain an understanding of the degree and nature of ion-solvent and ion-ion interactions (i.e. ion-pair formation) in this solvent, electrical conductivity, viscosity, density, and n.m.r. studies on solutions of metal difluorophos- 4 phates and related compounds have been carried out. Earlier work has described electrical conductivity studies on solutions of the.alkali metal difluorophosphates. Because of its importance to the present work and the fact that it is as yet unpublished and hence.not readily available, the results will also be given in detail here. (Table 2)
2.2 Experimental
A) Preparation and purification of materials.
1.- Difluorophosphoric acid
Commercial HPO2F2, supplied by Ozark-Mahoning Chemical
Company, Tulsa, Oklahoma, was purified by double distillation at a pressure of 100 mras. of mercury and a temperature range of
50-52° in the apparatus shown in Fig. 1. A more detailed Fig. 1. Difluorophosphoric Acid- Distillation Apparatus 10 description of the procedure involved in purifying the acid was 4 given previously. ii Metal difluorophosphates • *
The preparation of anhydrous alkali metal difluorophos• phates by the reaction of alkali metal chlorides with difluoro- phosphoric acid according to,
MCI + HP02F2 -* MP02F2 + HC1 (2.1) was reported previously.^" In the present work anhydrous calcium and barium difluorophosphates were prepared in an analogous manner by the reaction of the alkaline earth metal chlorides
(previously dried by heating to 190° for 18 hours in an oven and stored over phosphoric oxide in a vacuum desiccator) with the acid. The samples were stored over phosphoric oxide in a vacuum desiccator until used. Microanalyses were obtained from the
A. Bernhardt Microanalytical Laboratories, Germany and the results are given below:
Ca(P02F2)2 Ba(P02F2)2
calculated obtained calculated obtained
%? 25.60 25.43 16.26 16.39
%F 31.40 31.65 22.40 22.70
%Ba 40.4.6 40.35
B) Electrical Conductivity
The design of the cell used to measure the conductivities of solutions in difluorophosphoric acid is shown in Fig. 2. The cell could be attached to the distillation apparatus at K by means of the B19 ground glass cone L. The cell has two electrodes and Fig. 2 Electrical Conductivity Cell a cell constant of approximately 6. The capacity of the cell
was about 400 mis..
The cell was cleaned with aqua-regia and the electrodes
were plated with platinum blarcR* by electrolysing a chloro- platinic acid solution prepared according to Jones and Bollinger.
The solution consisted of a 0.3% solution of chloroplatinic acid
in 0.025N hydrochloric acid with 0.02$ lead acetate added. A
current of ten milliamps. was passed for fifteen minutes with a
reversal of current every ten seconds. The cell was steamed out,
dried and then calibrated using aqueous potassium chloride solu-
tion according to the method of Lind, Zwolenik and Fuoss.
The cell was replated and recalibrated periodically. All
measurements were made with the cell immersed in an oil bath regulated by means of a mercury-thallium thermo-regula'tor at
25 i.002°. The temperature of the oil bath was measured by
Beckmann thermometers which had been calibrated against a plati•
num resistance thermometer.
For the addition of solid material to the cell the ap•
paratus shown in Fig. 3 was used. It consisted of a 'T'-shaped
glass tube, with B 19 ground glass sockets at the ends 0 and P,
and a B 24 ground glass cone with an extension at M. The corks
at 0 and P were made of Teflon and they were tightly fitted with
stainless steel pistons A and B respectively. The flat 'runners'
which were also made from Teflon interlocked at Q and.lay'on the
bottom of the glass tubes. The compound to be added to the cell
was weighed into small, preweighed, dry glass boats which were
stored in the side arm. It was found that approximately eight
"to minimize polarisation effects. Fig. 3 Injector,, used for Solute Additions to the Conductivity Cell boats could be accommodated in the side arm. If the solute was thought to be hygroscopic the samples were weighed out in the dry box. However, to remove any water which may have been ab• sorbed in the weighing-out process the loaded injector was connected to a vacuum pump via liquid nitrogen traps, warmed and evacuated. The injector was detached from the pump after a few hours and stored until use in the dry box.
A conductivity run was carried out in the following man• ner: by means of a B 3-9 inner ground glass joint the conduc• tivity cell was attached to the distillation apparatus at K and flushed out with dry air. Difluorophosphoric acid was distilled directly into the cell and acid obtained in this way usually had —L —1 —1 a conductivity of about' 2.6 x 10ohms" cm7 * At all times in handling the acid great care was taken to exclude water. In a test run on the pure acid, a gradual increase in the specific conductivity occurred with time; over a period of seven hours a 0.5% increase in Ji. was observed'! Solutions for conductivity measurements were prepared as follows: difluorophosphoric acid was distilled directly into the cell which was weighed before and after addition of acid. The B 21+ stopper was removed and the injector was quickly inserted into the cell at F. Mercury • was poured into the glass tubes holding the platinum electrodes and then the cell and injector were placed on a stand in the oil bath. Addition of solute was achieved by pushing a glass boat from the side arm into the main tube by the piston A and then the boat was moved by piston B along N and. pushed into the acid. •^perhaps due to water impurity. , The cell was well shaken after each addition of solute to ensure good mixing and then returned to the oil bath. After sufficient time had elapsed to allow for temperature equili• brium (fifteen minutes) the resistance measurements were made.
The cell was then removed from the bath, reshaken, and the con• ductivity redetermined until no further change in resistance was observed. In this way variations in conductivity due to improper mixing were eliminated.
Liquid solutes were added to the cell by means of a 2ml. microburette (R.G.I. Inc.) shown in Fig. 4. A Teflon adaptor was constructed (inset Fig. 4) so that the microburette could be attached to the cell; it consisted of a Teflon sleeve with a
B 19 socket at one end and a Teflon nut containing a Viton '0* ring. The adaptor was held at position B on the glass barrel of the burette by placing the Teflon nut, A, containing the T0T ring at B and screwing up the sleeve into the nut until it was held firmly and did not slip down the glass tip. The dry micro• burette and liquid solute were placed in the dry box. The burette was rinsed out and then filled with the solute. The glass tip was wiped and then capped with a tube fitted with a
B 19 cone which fitted into the Teflon B 19 socket sleeve. The apparatus was removed and attached to the cell, the tube being removed at the last moment. Additions were made by screwing up the plunger and noting the initial and final volumes. By making use of the density of the solute the mass of the addition and thus the resulting molality of the solution could be determined.
Resistances of solutions were measured on a precision 16
Fig. k Microburette, used for Solute Additions to the Conductivity Cell
f> cr O SiiuiWLi
MM
B
in 17
a~c resistance bridge which has been previously described by 25
Daggett. A 2000c/s oscillator was employed as the source and
a telephone headset was used as the null detector. Throughout
this work, specific conductivity will be referred to by the symbol J{_ .. Some measurements were made at 1000 c/s and these gave satisfactory agreement with those at 2000 c/s. C) Nuclear magnetic resonance (n.m.r.)
A desired amount of solute was added to a weighed, dry
n.m.r. tube (Varian Associates, analytical n.m.r. sample tube,
part no. 905-370) of 0.5 cm.o.d. which was then reweighed and
transferred to a dry box. About 1ml. of HPO2F2 was added to the
tube which was again reweighed. The tube was once more re•
turned to the dry box where a capillary tube (0.15 cm.o.d.) of
HP02F2 used as an external reference was added. During these handling procedures the tube was tightly capped. Finally the n.m.r. tube was * flame sealed* and stored in liquid, nitrogen until the spectra were run (it usually took 4-5 days to ob• tain all the spectra). It was found that addition of the HPQ^Fj ex• ternal reference to a n.m.r. tube of HPO2F2 caused no new reson• ances to appear in the fluorine or phosphorus spectra.
The chemical shifts, relative to the pure solvent, were measured with a Varian HA 100 high-resolution spectrometer,• operating at 100.0 Mc/s for 1H, 94.07 for 19F and 40.43 for 31P.
Chemical shifts were measured by the * side-band* technique, in the case of the phosphorus and proton spectra and on calibrated, chart paper for the fluorine spectra. In the "^F spectra the region around each peak of the doublet was-examined by locking-in on the other peak of the doublet. 18
D) Viscosity-
Viscosity measurements of solutions of several metal di-
fluorophosphates in difluorophosphoric acid as solvent were
carried out at 25°. Solutions were made up in the drybox by add•
ing a weighed amount of HP02F2 to a known weight of the metal di-
fluorophosphate contained in a test tube which was then thorough•
ly shaken. In all cases the density and viscosity determina•
tions were made on the same solution. 26 A standard Ostwald viscometer was employed which was
calibrated using distilled water and dimethyl sulphoxide as
standards. The temperature was controlled by placing the visco•
meter in a 25-0.1° waterbath. To make sure that the drying
tubes which fitted on the openings of the viscometer and would
be needed for the HPO2F2 determinations would not affect the results several measurements were taken using the calibrating
liquids with the drying tubes in position as well as removed. No
difference in the results was observed.
The viscometer was first washed with concentrated nitric
acid and then several times with deionized water; it was then
dried and flushed-out with dry air for 1-2 hours. Ten mis. of
liquid were pipetted into the wide bore tube of the viscometer which was placed upright in the water bath and then left for an
hour to attain 25° - the tubes being capped to prevent evapora•
tion before measurements were made. Several runs were made on
each solution until consistent results were obtained. Great care was taken to make sure that no dust particles were trapped in
the capillary. In the case of the HP0oF solutions the viscometer was charged with the solution in the dry box.
E) Density
The clean, dry specific gravity bottles of 10 mis. capa•
city were weighed and then filled with liquid. The bottle was then suspended to its neck in the water bath and left for an hour. On further examination it was usually found that the liquid level was mid-way up the capillary of the stopper. The capillary was filled by means of a fine pipette. Excess liquid which expanded out through the capillary was mopped up by tissue paper. 'When it.appeared that the bottle and capillary were full of liquid and at 25° the bottle was removed and placed in a beaker of cold water. The bottle was then thoroughly dried and weighed. This procedure was repeated three times for each determination. To calibrate the bottles, liquids of accurately known density were used, namely mercury and water, care being taken to exclude all air bubbles. For the solutions in di- fluorophosphoric acid the weighed density bottle was filled in the dry box and then transferred to the water bath.
2.3 Results and discussion A) Electrical conductivity The results of the conductivity measurements on solutions of metal difluorophosphates in difluorophosphoric acid at 25° are given in Table 2. As all the solutions were made up by weight the concentrations are expressed in molal units, m. In each case a plot of the specific conductance, X. against molality was made (Fig. 5). As density and viscosity data are 20
TABLE 2
•SPECIFIC CONDUCTANCES OF SOME DIFLUOROPHOSPHATES AT 25°C
LiP02F2 KP0 F RbP02F2 9 2 2 10 dm 10 ^m 102m lO2^ ohm. .cm. ohm. cm. .ohm cm. ^"
0.000 2.482 0.000 2.410 0.000 2.499 0.306 2.673 0.158 2.557 6.629 2.662 1.020 3.464 0.566 2.740 2.051 3.313 2.190 4.528 1.366 3.153 4.056 4.314 4.429 6.379 2.435 3.889 7.084 5.325 6.773 8.086 4.043 4.925 11.22 7.329 9.733 10.15 6.054 5.939 I6.84 10.56 13.91 12.73 9.310 7.658 23.63 13.81 19.69 16.59 13.33 10.10 31.88 17.72 27.66 20.46 18.53 12.79
37.20 24.43 23.88 15.21 CsP02F2 49.34 29.18 29.65 17.77 10 2m 104U 60.08 32.21 34.37 19.82 70.55 34.56 40.28 22.46 ohm. cm. 48.26 24.27 0.000 2.443 O.468 2.565 1.409 2.923 NaP02F2 NH^P02F2 3.191 3.708 102m lO1^ 102m 10^ 6.155 5.041 ohm;"Icm.~1 ohm. -lcm.-l 9.751 6.665 13.48 8.395 0.000 2.503 0.000 2.473 17.92 10.51 0.222 2.557 1.977 3.546 24.29 13.64 0.761 2.868 5.864 5.351 (P0 2.859 4.500 11.44 9.110 Ca 2F2)2 18.75 2 -_J1 6.003 6.794 13.35 10 m lcAtt ohm - cm 8.220 8.266 27.15 13.71 11.16 10.12 35.60 23.91 0.000 2.553 15.13 12.26 1.120 2.758 (P0 20.95 15.15 Ba 2F2)2 2.470 4.628 25.36 17.96 4.327 5-793 2 31.37 20.67 10 m 10^ 7.533 8.676 22.81 -1 36.33 ohm.~^cm. 10.79 10.30 42.46 24.75 14.83 11.77 47.79 26.40 0.000 2.823 19.66 13.53 1.736 3.783 23.50 14.27 6.033 8.075 7.646 9.800 10.08 12.39 13.44 15.29 16.02 17.12 20.08 20.21 24.23 22.70
22
available for solutions of LiP02F2, KP02F2, RbP02F2 and Ba(P02F2)2
(see later), equivalent conductivities could be calculated for
these solutes and these are given in Table 3, and in Fig. 6 plots
of equivalent conductance versus the square root of ionic strength
are shown. The equivalent conductance values, A , were calcula•
ted from the expression, TV = lOQoXi c • '
where C is the concentration in gram equivalents per litre. The
value of Xi, the specific conductivity caused by the addition of
solute, was calculated fromXi = X. —X,o, the observed conduc•
tance less the conductance value at infinite dilution1'! For the
alkali metal dif luorophosphates K.o was taken as 2.25 x 10~^'ohm~"'"cm7^
but for the calcium and barium salts values for l^o were
2.15 and 1.90 x 10"^ ohm~x cm, . These equivalent conducti•
vity values may be corrected for the change in the bulk vis-
cosity of the solution by multiplying by Vj /yj } where Yj is o the viscosity of the pure solvent and Yj the viscosity of the
solution.2 7 The Arj
values are also given in Table 3; the effect of the viscosity correction causes a lowering of the
equivalent conductivity values, which is greatest for the more
concentrated solutions. In the case of Ca(P02F2)2 solutions
neither density nor viscosity measurements were made, thus no
viscosity correction can be applied and the equivalent conduc•
tivities (cf. Table 3) were calculated assuming the same value
for the solutions as that of the pure solvent, 1.5$3g./ml.
The most striking feature of these results is that the conductivity values are considerably smaller than those obtained TABLE 3
EQUIVALENT CONDUCTANCES OF SOME ELECTROLYTE'S AT 25°
LiP02F2
2 io c ioc^ 10*00., A ^Vlc ohm." era."1 * O.484I • 0.6958 0.423 8.739 8.687 1.613 1.270 1.214 7.526 7.383 3.459 1.859 2.278 6.585 6.269 6.984 2.643 4.129 5.912 5.404 10.66 3.265 5.836 5-475 4.851 15.29 3.910 7.90 . 5.167 4.438 21.78 4.667 10.48 4-811 4-070 30.69 5-540 14.34 4-671 3.910 42.89 6.549 18.21 4.245 3-553 57-24 7-566 22.23 3-683 3-240 76.00 8.718 26.93 3.542 2.940
KP02F2
2 10C2 10 C 10S. -1 A ohm cm. 0.2500 0.5000 0.307 12.26 12.27 0.6953 0.9462 0.490 5.474 5.458 2.112 1-454 0.903 4.276 4.233 3.925 1.961 1.639 4-175 4.092 6.379 2.526 2.675 4-193 4.046 9.537 3.086 3.669 3.668 3.655 14.63 3-625 5.406 3-697 3.390 21.66 4.654 7.65 3.624 3-182 28.93 5-379 10.54 3-643 3-071 37.13 6.093 12.96 3.490 2.855 45-91 6.776 15.52 3-360 2.643 53-04 7.283 17.57 3-312 2.524 61.92 7.669 20.21 3.268 2.445 73-77 6.569 22.02 2.988 2.384
RbP02F2
102C 10C2 Av ohm._J-cm. --1 . 1.011 1.005 0.412 4.075 3.976 3.241 1.800 1.068 3-295 3.193 6.399 2.530 2.064 3.225 3.036 11.14 3.336 3.575 3.209 2.880 17-60 4.195 5.579 3.170 2.729 26.29 5.128 8.31 3.161 2.665 36.71 6.059 11.56 3.149 2.623 49.22 7.015 15.47 3.143 2.595 TABLE 3 (cont'd)
Ba(P02F2)2 102C 10C~2 1043il A ohm."cm."1 2.738 2.867 1.883 3.436 9.453 5.326 6.175 3.265 11.94 5.986 7.900 3-308 15-67 6.857 10.49 3-347 20.78 7-900 13-39 3-221 24.67 8.603 15-22 3-084 30.72 9-601 18.31 2.980 36.83 10.51 20.80 2.823
Ca(P02F2)2
1.768 2.304 1.719
111?? W5 16.65 7.069 22.67 8.247 29.71 9.442 35.20 10.28 25
Fig. 6 Equivalent Conductivities of Some Difluorophos• phates at 25° plotted against the Square Root of the Ionic Strength for strong bases in other associated protonic solvents. Table
4 compares values of specific conductances at 0.1 molal concen•
tration for strong bases in various solvents.
TABLE 4
Specific Conductances of Some Potassium Salts in Various Solvent
Solvent Solute Specific conductance of 0.1m solution
Z 4 HP02F2 " KP02F2 7.81 xlO~
HSO^F 10 KSO^F 220 xlO-4
xlO"^ H2S04 19 KHSO 4 200 28 H20 KOH 275 xlO-4-
The lower conductivity values observed for the HP02F2 solutions
are either due to ionic mobilities in HP02F2 being much less
than those in the other solvents listed in the Table or else due
to the metal difluorophosphates being incompletely dissociated
electrolytes in HPOgFg.
Since the dielectric constant of HPOgFg is
not knovm, making it impossible to predict the importance of ion
pairing in this solvent, the results will be examined on the
basis of two models: (i) ion-pairing is not important in this
solvent and the metal difluorophosphates are fully dissociated
electrolytes, (ii) ion-pairing is important and the metal di-
fluorophosphates are incompletely dissociated into free ions.
Model (i)
In solvents where the mobilities of the autoprotolysis ions are much greater than the mobilities of other ions (because of a proton transfer mechanism of conduction for the former ions) strong bases exhibit almost identical specific conductivity curves at low concentrations with small deviations noticeable only in more concentration solutions. As the conductivity curves for the alkali metal and alkaline earth metal difluorophos phates deviate from each other at even the lowest concentrations measurable, it must be concluded that, if the difluorophosphates are fully dissociated, the PO2F2*" ion does not show abnormal conduction. This conclusion is consistent with the low conduc• tivities observed for these solutions but is surprising in view of the associated nature of the solvent.
The conductivity of the alkali metal difluorophosphates decreases in the order Li)>Na^ NH^} Rb^> Cs at any given con• centration. As each solute has the common ion, PO2F2", then the differences in the conductivity must be due to differences in the mobilities of the cations if it is assumed that complete dissociation of these salts occurs. This order of cation mobility is opposite to that found by Gillespie et al. in their 10 29 conductivity measurements in HSO^F and HgSO^ solutions where they found the order Cs^> Rb)> Nh\^
also completely dissociated it follows that the mobility of the
Ca ion is less than the mobility of the Ba ^ ion i.e. the
smaller ion has the smaller mobility, a reversal of the situa•
tion noted above for the alkali metal cations.
Model (ii)
If it is assumed that ion-pairing is important in HPO2F2 and the metal difluorophosphates are only weakly dissociated in this solvent, then the relative orders of the conductivity values
of these salts may reflect their relative degrees of dissociation. 30
In water } where ion-pairing is general though not extensive,
for the alkali metal salts of the oxy-acids the metals of higher
atomic number show more ion-pairing. Thus cesium salts tend to
exhibit more ion-pairing than lithium salts and consequently
produce fewer ions in solution. If this tendency of ion-pairing were to occur in HPO2F2 then lithium difluorophosphate would be.
the most dissociated and would be expected to exhibit the highest
conductance of the alkali metal salts, as is observed.
The specific conductivity curves for calcium and barium
difluorophosphates (Fig. 5) show considerable curvature with
Ca^Ba and at similar concentrations barium shows a greater con•
ductance than calcium. If it is assumed that these conductivity
curves simply reflect the differences in degrees of dissociation
of these solutes then clearly Ba(P02F2)2 is the most dissociated.
In the previous discussion of the possibility of ion-pairing in
the alkali metal salts it would seem that for the Group II metals with the increase in cationic charge, the formation of ion-pairs in quite dilute solutions would be the rule rather than the
30 exception. Davies reports that in water, for salts of the in•
organic oxy-acids, ion-pairing increases with atomic number and ionic radius of the cation, as for the Group I metals. This
shows that-ion-association cannot be explained by a purely
electrostatic theory in which the solvent functions only as a medium of uniform dielectric, for if so, the smallest ions would show most ion-pairing. However, Davies also reports that for .the hydroxides, fluorides and carboxylates of the alkaline earth metals the opposite order prevails, the larger barium ion giving the least ion-pairing. This order is consistent with the results given here in that the barium salt shows greater conductivity than the calcium salt at the same concentration and so is pre• sumably more dissociated.
In view of the fact that the specific conductivity of
HPO2F2 is high (in fact of the same order of magnitude as the specific conductivities of dilute solutions of the metal di- fluorophosphates) and as it is not known whether solvent self- dissociation or impurities are giving rise to this conductivity, quantitative treatment of the conductivity data is difficult.
Lack of information on the dielectric constant of HPC^Fg adds to the difficulty of treating th£sedata. When the specific con• ductivity curves are extrapolated to infinite dilution it is noted that the curves do not pass through the origin nor through the point corresponding to initial solvent conductivity. If it is assumed that the extrapolated conductivity values are due to impurities in the solvent and the difference between these values and the original solvent conductivity values are due to repression of solvent autoprotolysis then conductivity values
— Xx> may be taken as reasonable values for the conduc• tivity of added solute. Using these IKi*-values equivalent con• ductivity values for the metal difluorophosphates have been cal• culated and are plotted against the square root of the ionic strength in Fig. 6. Qualitatively it can be seen that at high concentrations the equivalent conductivities converge for the various solutes while at low values of ionic strength the curves diverge with a greater increase in «/V on dilution for L/LPO2F2 consistent with this solute showing the greatest degree of dis• sociation. Moreover at very high concentrations CofPOgFg^ shows the greatest decrease in W consistent with it being the least dissociated.
The interpretation of the conductivity data on the basis of model (ii) appears to be the more reasonable of the two model proposed i.e. ion association in HPO2F2 is important and metal difluorophosphates are weakly dissociated in this solvent.
Further evidence supporting this conclusion was obtained from th n.m.r., viscosity and density studies to be discussed later in this chapter and also from the fact that tetraphenylarsonium chloride gives more highly conducting solutions than the metal difluorophosphates (to be presented in Chapter III).
B) Nuclear magnetic resonance
The ^-H, 19F, and chemical shifts of the solvent peak relative to an external reference of HPOoFp were measured for a number of solutions of alkali metal salts in HPO^Fg. The results are given in Table 5 and the uncorrected values are plotted against the molal concentration in Figs. 7-9. The chemical shifts were measured for solutions over the same concentration range as that used in conductivity measurements and as a result the shifts are very small and correspondingly very sensitive to impurities such as water and particularly paramagnetic species.
In spite of these difficulties reasonably smooth curves x^ere ob• tained for all the solutes studied except the resonance of
KP02F2 solutions. This salt showed the smallest dependence of
"Hi chemical shift on concentration and gave results which were very scattered and difficult to reproduce. Chemical shifts are observed for all three nuclei when salt is added and the nature of the shift depends on the alkali metal.
The shifts were corrected for bulk diamagnetic suscep• tibility effects according to the equation1*',
S = H-H ref. + 2_TT (Xw^ef. —>X ir) H ref. 3 v
The susceptibilities of the solutions x^ere calculated from Pascal constants and the Wiedmann mixture law. The molar diamagnetic susceptibility of difluorophosphoric acid was taken as
38 x 10 ^c.gs.; this value was estimated by using that of ortho- phosphoric acid (43.8 x 10~^c.gs.)^ and substituting two fluorine
1 9.1 x 10""6c,gS.) for two hydroxyl groups (12.0 x 10""^c.gs.) .•^
Corrections for the bulk diamagnetic susceptibility effects were estimated and are given in Table 5. The application of these cor rections does not alter the nature of the results obtained; the TABLE 5
1<7 31 ^•H, F and P Chemical Shifts* for Solutions of MP02F2 in HP02F2
2 10 m. Chemical Shift (p.p.m.) Suscept. corr (p.p.m.) H 19W 31T
LiP02F2
6.792 +0.0142(.0010) -0.0627(.0002) -0.0546(.0049) +0.001 17.06 +0.0418(.0030) -0.133 (.001) -O.I46 (.005) +0.004 27.90 +0.0620(.0031) -0.191 (.001) -0.207 (.017) +0.008 32.59 +0.0756(.0025) -0.224 (.002) -0.259 (.006) +0.010 50.44 +0.0908(.0021) -0.311 (.001) -0.383 (.005) +0.016
NaP02F2
6.503 +0.0175(.0006) -0.0788(.OO64) -0.0776(.0074) +0.003 12.2,7 +0.0240( .0018) -0.138 (.007) -0.165 (.010) +0.006 28.39 +O.O415 (.0030) -0.288 (.004) -0.413 (.012) +0.014 31.85 +0.0479(.0020) -0.316 (.004) -0.447 (.015) +0.017 47.79 +0.0559(.0020) -0.442 (.004) -0.670 ).017) +0.025
KP02F2
9.030 +0.0282( .0010) -0.0839(.0053) -0.124 (.008 ) +0.007 17.42 -0.019 ( .002) -O.I84 (.003) -0.236 (.003 ) +0.013 25.62 +0.026 ( .002) -0.207 (.004) -0.294 ( .011) +0.018 35.84 -0.029 ( .002) -0.369 (.005) -0.469 (.005 ) +0.029 45.42 -0.455 (.005) -0.531 (.015 ) +O.O35
NH4P02F2
8.633 -0.0144(0012) -0.0999(.0003) -0.0588( .0015) +0.008 14.89 -0.0144( .0006) -0.156 (.003) -0.114 (.0005 ) +0.014 32.73 -0.0679( .0002) -0.351 (.005) -0.277 (.00 8 +0.033 34.74 -0.0725( .0002) -O.366 (.004) -O.304 (.002 ) +0.035 56.03 -O.132 ( .001) -0.553 (.010) -0.LS2 (.013
RbP02F2
5'.040 -0.0667(.0007) -0.0596( .0035 +0.004 15.50 -0.0276( .0023) -O.I9O (.003) +0.014 21.97 -0.0432(.0005) -0.271 (.004) -0.269 (.005 ) +0.022 32.47 -0.0790( .0021) -0.391 (.005) -0.410 (.008 ) +0.032 37.74 -0.0946( .0030 -0.443 (.005) -O.472 (.011 ) +0.038 33
TABLE 5 (cont'd)
2 10 in Chemical Shift (p.p.m.) Suscept. corr.
CsP02F2
11.89 -0.0132(.0004) -0.152(.003) -0.157(.004) +0.017 14.70 -0.0253(.0012) -0.190(.002) -0.180(.003) +0.021 20.75 -0.0781(.0010) -0.273(.004) -0.259(.011) +0.030 32.62 -0.127 (.002) -0.428(.005) -0.408(.010) +O.O49 43.86 -O.I64 (.006) -0.517(.005) +0.066
* Values for the shifts are the average values of at least two • determinations; the values in the brackets are the average deviations.
36 corrections simply shift.the curves slightly to higher fields and so for discussion purposes the uncorrected curves, shown in
Figs. 7-9 only need be considered.
The effect on the solvent proton chemical shift of adding salts has been investigated in a number of solvents, notably
2 H20^ , H2S04 -^and HSO^F"^. In a series of compounds studied by
Schneider, Bernstein and Pople-^ it was found that in all cases association through hydrogen bonding shifts the proton signal to low field. Association through hydrogen bonding will thus pro• duce a low-field shift whereas high-field shifts caused by sol- vated ions can be attributed to there being fewer hydrogen bonds between the solvation layer and the bulk solvent than between the solvent molecules themselves. This break-up of solvent structure has been examined extensively in water by proton mag- • netic resonance-^2. Hindman^D divided the effects of ions on the proton magnetic shielding by electrons in water molecules to high- and low-field shifts. The high-field shifts were at• tributed to the breaking of hydrogen bonds in the process of re• orientating the water molecules by the ion, related to the forma• tion of the primary solvation layer, and also the breaking of additional bonds beyond the primary hydration layer i.e. in the in the structure broken region. If an: ion were capable of in• ducing more hydrogen bonding in the solution than existed in the pure solvent a low-field shift would occur.
Before proceeding with' a discussion of the results ob• tained in this work, it is convenient to consider the results of analogous 1H n.m.r. studies by other workers on solutions of 38
alkali metal hydrogen sulphates in sulphuric acid and alkali metal fluorosulphates in fluorosulphuric acid^*". In both cases it was observed that addition of salt caused the solvent proton signal to shift to low-fields. The extent of this shift de• pended on the concentration of the salt and the alkali metal cation involved — the largest low-field shift being produced by the largest cation. These results were explained as follows:
The anion (SCUF~ in the case of HSO-F solutions and HSO, ~ in } 3 4
S the case of H2 °4 solutions) interacts with the solvent molecules
in such a -way as to increase the number of hydrogen bonds present
in solution over that occurring in the pure solvent and hence
causes a shift to low-field. Superimposed on this is the effect
of the cation which by virtue of its solvation breaks up the
solvent structure and so decreases the number of hydrogen bonds-
present and thus gives rise to a high-field shift in the proton resonance. The smaller the cation the greater is its degree of
37
solvation and the greater is the up-field shift caused by it. The overall downfield shift observed for all the salts is due to the fact that the effect of the anion is considerably greater than that of the cation. The results of the present studies on alkali metal di- fluorophosphates dissolved in HPO2F2 are to be contrasted with those described above. Lithium and sodium difluorophosphate actually cause a shift to high-field in the proton resonance of the solvent while rubidium and cesium difluorophosphate cause a shift to low-field. It would appear that unlike the situation with K^SO^, and HSO^F the down-field shift caused by the anion is 39
of comparable magnitude to the up-field shift caused by the cations.
It is important to note, however, that as in the previous systems studied, HSO^F and H^SO^, the smaller the cation the greater is the shift to high-field, suggesting that in HPO^Fg the cations are solvated, the smaller ones being more solvated than the larger ones. The. problem which remains is discovering why the
P(02^2~ ion does not cause a large solvent proton shift to low- fields as do the HS0^~ and SO^F" ions in their respective sol• vents. The answer probably lies in the fact that the metal di- fluorophosphates are extensively ion-paired in HPO^F^. The alkali metal cations may be pictured as being solvated by sol• vent molecules with difluorophosphate anions present within the cation solvation shell. Hence the cations by virtue of their solvation cause an up-field shift in the ~*"H resonance while the
P02F2"" ions have less effect on the solvent chemical shift than they would if they were free ions as is the case for HSO^- ions
in HoS0, and SO_F~ ions in HSCLF.
When alkali metal difluorophosphates are added to HPO2F2 a single ^9p resonance is observed (actually a doublet due to coupling with phosphorusf^which is an average signal due to fluorine atoms in the solvent molecules and fluorine atoms in the PO2F2" ions. The shift of this signal with respect to pure
HPOgFg was measured for a number of solutions and the results are presented graphically in Fig. 8. Due to the fact that POgFg" in is an anion it would be expected that the 7F resonance would occur at higher field strengths than the "^p resonance of HPO2F2
and. hence the addition of MP02F2 to HPOgFg might be expected to 40
cause a F upfield shift relative to pure HPO2F2. However, as can be seen in Fig. 8, all of the shifts are to low-field. In analogous studies on solutions in fluorosulphuric acid it was found that the 19F resonance is indeed shifted to high field as metal fluorosulphates are added to the solvent. Again it is concluded that the expected high-field shift due to added POgFg" ions to difluorophosphoric acid is not observed because the
^®2^2~ i°ns are not 'free* but are present in the solvation sphere of the cation. The down-field shifts which are observed are presumably due to the interaction of the cations with the solvent molecules. It is interesting to note that the largest down-field shift is observed for the largest cation, Cs+, v' which has the most solvent molecules in the first solvation sphere and hence in the immediate environment of the cation. 31 The P n.m.r. studies of solutions of alkali metal di- fluorophosphates in HPO2F2 show, as was the case for the 19F spectra, that added metal difluorophosphate shifts the "average" resonance to low-fields. Again, the expected high-field shift
P0 F ions is not due to the added 2 2~ observed and the dov/nfield shifts are presumably due to the interaction of the solvent mole- 1 19 cules with the cations. Unlike the H and F n.m.r. results there is no correlation between the magnitude of the observed shift and the size of the alkali metal cation although Li+ does seem to show a significantly smaller downfield shift than the other cations and this is probably due to the fact that there are fewer solvent molecules in its first solvation sphere. It is to be expected that the cation-solvent interactions will have less 41 effect on the shielding constant of the P nucleus than on the
H or F nuclei since the phosphorus atom is positioned at the centre of the molecule, screened by two fluorine atoms and two oxygen atoms from the cation. This probably accounts for the fact that little difference is observed in the plots in Fig. 9 for the Na, K, Rb and Cs salts.
2.3 C) Densities In Table 6 the densities, d, of some alkali metal di- fluorophosphate solutions in difluorophosphoric acid at 25 are given and in Fig. 10 the densities of these solutions are plotted against concentration, m, expressed in moles per lOOOg. of solvent. It appears that inorganic cations cause an increase
centration. Gillespie and Wasifp found a similar trend for metal bisulphates in sulphuric acid but a decrease in density with increasing concentration was observed for organic cations. Density changes can be more conveniently discussed in terms of the apparent molar volume of the solutes rather than in terms of the densities themselves. The molar volume, 0, of an electro-
3 lyte is related to the density of the solution, d, by the equation,
the pure solvent and m, the molality of the solution. The apparent molar volumes given in Table 6b v;ere cal• culated from the equation. It can be seen that there is no marked change in the apparent molar volume with concentration and TABLE 6 o a) Densities and Viscosities of Some Solutes in HPO2F2 at 25
LiP02F2 102m D^5 ^ 0 0.000 1.5827 (.0001) 5.200 (.040) 7.185 1.5855 (.0002) 4.544 (.040) 48.9 18.29 1.5911 (.0003) 4.392 (.018) 48.2 23.36 1.5965 (.0002) 4.330 (.055) 48.7 32.30 1.5984 (.0005) 4.350 (.075) 47.8
RbP02F2 " 0.000 5.200 (.040) 10.06 1.6003 (.0004) 4.536 (.030) 45.0 22.07 1.6197 (.0005) 4.366 (.090) 48.0 35.45 1.6413 (.0003) 4.310 (.193) 48.0 KP02F2 0.000 1.5824 (.0002) 5.769 (.001) 12.60 1.5960 (.0004) 5.122 (.020) 44-0 23.54 I.6O64 (.0001) 5.029 (.018) 47.0 39.16 1.6231 (.0003) 4.306 (.021) 45.6 54.49 I.638O (.0008) 4.617 (.040) 45.3
Ba(P02F2)2 0.000 5.722 (.020) 13.57 5.011 (.131) 23.05 1.6406 (.0008) 5.269 (.065) HO 29.95 1.6599 (.0005) 5.407 (.055) 106
(CH0),NC1 3 4 0.000 5.658 (.052) 16.45 3.963 (.025) 27.34 1.5745 (.0011) 3.724 (.035) 49.52 1.5699 (.0012) 2.900 (.015) TABLE 6 (cont'd)
b)
Solute Mean 0 Mean 0^ Mean mis. mis. $H2S0k
LiP02F2 48 tl -16 tl -7.0 -7.0
KP0 F 45 -19 t2 -1.0 +1.7 2 2 -2
RbP02F2 47 ±2 -17 -2 +5.0 +7.7
Ba(P02F2)2 108 t2 -20 t2 -12 -24.3
c) viscosity depression,AT\ 0 • molality, 10Si. Solute 0.100 0.200 0.300 0.400 LiPOj^ -0.707 -0.804 -0.847 -0.872
KP02F2 -0.519 -0.888 -1.27 -I.46
RbP02F2 -0.655 -0.820 -0.869 -0.890
Ba(P02F2)2 -0.637 -0.642 -0.222
(CHjNCl -1.09 -1.83 -2.29 -2.56 J 4
45
also the apparent molar volumes are virtually the same for all of the alkali metal salts studied. This is to be contrasted
with the results obtained for solutions in H SO.^ and Ho0^"^ 2 if 2 where the apparent molar volume is a function of the concentra•
tion and nature of the salt added. In the water and HgSO^ work
values of 0*, the apparent molar volume of the cations, were
calculated assuming that the partial molar volumes of the anions
are the same as the molar volumes of the solvent molecules.
If it is assumed that the apparent molar volume occupied
by a difluorophosphate ion in solution is the same as that of a
difluorophosphoric acid molecule, then 0^ values may be obtained.
The values given in column 3 of Table 6b were calculated by using
a figure of 64 mis., the molar volume of HPO2F2, for the apparent molar volume of the difluorophosphate ion. Thus the apparent molar volume of the cation, 0*, is given by: 0 - 64 = 0*.
These are compared with the apparent molar volumes in aqueous
solution and in sulphuric acid and as can be seen the values are
considerably greater in the case of HPO2F2 solutions. It seems
likely that the 0^ values in difluorophosphoric acid are in•
correct. If a partial molar volume of less than 64 is assumed
for the dif luorophosphate anion then the values of 0"*" will be in
better agreement with those obtained in the other systems. In•
deed, if the difluorophosphate anions are present in the solvation
spheres of the cations and not 'free', as it is strongly suspec•
ted, then it is not unreasonable to expect that the anion will
have a smaller apparent molar volume than the solvent molecules. 19 32 It has been shown that in sulphuric acid • and water 46 the extent of solvation of metal cations decreases in the order
Ba^Li^ K^Rb. From Table 6b it can be seen that this same order prevails with increasing negative values of i.e. Ba2+ is most negative and Rb+ most positive. The values for the difluoro• phosphates are approximately the same and thus no information regarding the relative order of solvation numbers can be ob• tained as was possible in the solvents where the salts were fully dissociated.
2.3 D) Viscosities In Table 6a the viscosities for various solutions are given. The effect of the solute on the viscosity is given by,
" ^lo wnere Is tne viscosity of the solution. Interpolated values are given in Table 6c for the viscosity depression. All the solutes studied caused a decrease in the viscosity with tetramethylammonium chloride causing the largest depression. 38 In sulphuric acid it was found that metal hydrogen sulphates generally cause an increase in the viscosity, the ef• fect of the alkaline earths being particularly great, whereas organic molecules were found to cause a small decrease in the viscosity. Gillespie explained changes in the viscosity in terms of the effects of the cations on the structure of the solvent.
The small inorganic cations were considered to cause a 'tighten• ing1 of the structure of the solvent around the ion, pulling the solvent molecules strongly together and considerably restrict• ing their freedom of movement and thus increasing the viscosity.
However, in the case of the .large organic cations it was suggested that in addition to the effect caused by their solvation the bulky 47 groups would tend to disrupt the structure of the surrounding solvent, as they would not fit easily into it and this would probably be accompanied by a decrease in viscosity.
The viscosities of all the difluorophosphoric acid solutions are less than the viscosity of pure HPO2F2, thus indicat• ing the behaviour of large, bulky solutes and not of fully dis• sociated salts. It may be that the ion pair itself is structure breaking and would not easily fit into the structure of the sur• rounding solvent. Viscosity and density measurements were also made on tetramethylammonium chloride solutions in HPO2F2 as a comparison study. CHAPTER III
Miscellaneous Bases
A) Organic Bases
3.1 Introduction
A large number of organic compounds are bases in the 20 21 hydrogen fluoride and sulphuric acid ' solvent systems. Most substances containing oxygen, nitrogen or sulphur atoms that are not coordinatively saturated offer a lone electron pair capable of binding protons in acidic solvents. In HgSO^ amines are strong bases and are simply converted to their conjugate acids19. Nitrocompounds, however, being weaker bases, exhibit a variety of base strengths (m-nitrotoluene> nitrobenzene y p-nitrochlorobenzene^ 2,4-dinitrotoluene) and are not fully protonated. It was of interest to see whether a span of base strengths of such organic molecules can be observed in HPO^Fg solutions.
3.2 Experimental
A) Electrical conductivity studies were carried out at
25° on solutions of various organic bases in the cell shown in
Fig.. 2. The procedure adopted was the same as that outlined previously; solids were added by means of the 'injector1 and liquids by the microburette.
B) Preparation and purification of materials
Nitrobenzene: 'Analar* nitrobenzene was purified by fractional freezing followed by double-distillation under vacuo.
Benzoic acid: 'Analar' grade benzoic acid was recrys- tallized from water and dried over phosphoric oxide in a vacuum 49 desiccator. Oxalic acid: lAnalar* grade oxalic acid was dried in an oven at 150° and stored over phosphoric oxide in a vacuum desic• cator.
Reagent grade 2,4-dinitrotoluene, p-nitrotoluene, m-nitroaniline, p-nitroaniline, p-nitrochlorobenzene and p- dinitrobenzene were recrystallized from methanol, dried and stored over phosphoric oxide in a vacuum desiccator. 3.3 Results and discussion The results of the conductivity measurements on solutions of p-nitroaniline, m-nitroaniline, 3,4-dichloroaniline, 2,4-dinitrotoluene, p-nitrotoluene, p-nitrochlorobenzene, p-dinitrobenzene, nitrobenzene, benzoic acid and oxalic acid are given in Table 7 and are plotted against the molality of the solution in Fig. 11. It was found that the anilines and benzoic acid were completely soluble over the concentration range studied whereas oxalic acid, p-dinitrobenzene and 2,4-dinitrotoluene ap• peared virtually insoluble and caused essentially no change in
the conductivity when added to HP02F2. Assuming that the point at which further addition of solute causes no noticeable increase in the conductivity of the solution corresponds at least approxi• mately to the solubility of the solute it would appear that the compounds p-nitrotoluene, p-nitrochlorobenzene and nitrobenzene have solubilities 0.08, 0.034 and 0.15 moles/kg. respectively. The three anilines have nearly linear conductivity- molality curves which are similar, particularly at low concentra• tions, to the curve obtained for cesium difluorophosphate. This TABLE 7
Conductivities of Some Organic Bases in Difluorophosphoric Acid at 25°
p-nitroaniline m-nitroanlline 3,4-dichloroaniline
102m. 104X' 102m. 10*X 102m. 10^ ohrrT+cm. ohm_1crn. 1 ohm^cm -1
0.0000 2.435 0.0000 2.648 0.0000 2.569 0.2666 2.489 0.8946 2.538 0.4074 2.483 1.218 2.836 3.256 3.456 1.156 2.596 2.576 3.457 7.252 5.273 2.843 3.227 4.562 4.366 12.34 7.467 5-493 4.339 7.304 5.583 19.60 10:73. 8.738 5.630 10.30 6.849 27.41 13.63 12.50 7.039 14.51 8.542 35.85 17.58 16.97 8.635 18.56 10.11 43.64 19.78 21.04 10.03 22.87 12.67 50.75 21.53 25.30 11.62
2,4-dinitrotoluene p-nitrotoluene p-nitrochlorobenzi 0.0000 2.417 0.0000 2.458 0.0000 2.617 0.7450 2.476 0.8714 2.688 0.5345 2.699 1.932 2.474 2.377 3.325 1.711 2.850 4.267 2.456 4.638 4.233 3.380 2.892 insoluble 7.969 4.896 6.O64 2.804 12.88. . 4.851 insoluble insoluble
nitrobenzene p-dinitrobenzene benzoic acid 0.0000 2.893 0.0000 2.458 0.2224 2.962 0.0000 2.885 0.4769 2.602 0.5004 3.057 0.6382 2.940 1.406 2.895 0.9823 3.228 4.565 2.945 3.207 3.440 1.909 3.534 insoluble 5.947 4.186 3.077 3.912 9.825 5.101 4.374 4.331 oxalic acid 14.05 5.966 6.635 4-982 19.82 6.993 9.434 5.752 Q.0000 2.690 24.40 7.682 13.12 6.679 insoluble 27.10 8.074 16.83 6.845 insoluble Fig. 11 Specific Conductivities of Some Organic Bases in HPO^F^ at 25°
• r
© CsP02F2 A p-nitroaniline • m-nitroaniline O 3 , h , -dichloroanilin^ H nitrobenzene O benzoic acid
0 io
io2 x molality, m 52
suggests they are fully protonated but, like the metal difluoro- p'hosphates, are also incompletely dissociated. It is not sur• prising that the degrees of dissociation for the anilines are similar in view of a general similarity,in the sizes of the cations
of the ion pairs, viz. RM^'PC^Fp-.
Gillespie et al.1^ have compared the ionization constants
for some nitro compounds in fluorosulphuric and sulphuric acids
and due to the greater acidity of HSO^F organic bases are in
general protonated to a greater extent in this solvent. It -was
found that the ionization constants decreased in the order
nitrobenzene^ p-nitrochlorobenzene/> 2,4~dinitrotoluene; benzoic
acid was fully ionized in HSO^F and H^SO^. It would be expected,
therefore, as difluorophosphoric acid is known to be a weaker
2? acid than H^SO^ that the extent of dissociation of organic
solutes will be much less. The nitro compounds were in fact
found to be less soluble in HPO2F2 than the amines,with . the -
strongest bases nitrobenzene and p-nitrotoluene , showing the
greatest solubility. This lack in solubility on the part of the
nitro compounds is consistent with them being weaker bases than
the amines. The basicity of an organic solute relative to the
solvent must be great enough to at least cause hydrogen bonding with the solvent, otherwise it is unlikely to be able to suf•
ficiently disrupt the structure of the solvent to enable it to
dissolve.
The formation of the conjugate acid from an organic base
can be considered as a two-step process: 53
, B + HP02F2 ^ HB'T 02F2~ (3.1)
Jr + HB P02F2- ^ HB + P02F2~ (3.2) the first step being the protonation of the base to form an ion-pair and the second step being the dissociation of this species into free ions. It is probable that,in the case of benzoic acid, the lower conductivities relative to those observed for the organic amines are due to incomplete protonation, i.e. the benzoic acid is a weaker base than the amines and hence step (3.1) does not go to completion.
B) Inorganic Solutes
3.4 Introduction
It is well known that in sulphuric acid"^ the alkali and some other metal hydrogen sulphates behave as strong bases, analogous to hydroxides in water. Many salts of other familiar inorganic acids undergo complete solvolysis as is illustrated in the following examples,
+ KNO3 + H2S0^ - K + HSO^' + HNO3
+ NH^CIO^ + H2S04 = NH4 + HSO^" + HCIO^
As the alkali and alkaline earth metal difluorophosphates were prepared from the corresponding chlorides it was of interest to discover whether complete solvolysis of the chlorides takes place in solution. In an effort to examine the general usefulness of electrical conductivity studies in determining modes of re• action of solutes in difluorophosphoric acid, conductivity measurements were made of solutions of potassium nitrate and 54
carbonate in ll?O^F^.
3.5 Experimental
Electrical conductivity studies were carried out at 25° on solutions of various inorganic solutes in the cell shown in
Fig. 2. The solutes were added by means of the "injector" as described previously.
Tetraphenylarsonium chloride: commercial tetraphenyl- arsonium chloride obtained from Aldrich Chemical Co. was dried in. a drying pistol containing phosphoric oxide at 100° and 15 mms. pressure.
Tetramethylammonium chloride: commercial grade obtained from Matheson Co. Inc. was recrystallized from ethanol and dried by heating under vacuo.
The alkali earth and alkali metal chlorides were dried by heating to 190° in an oven' for 18 hours and then stored over phosphoric.oxide in a vacuum desiccator.
3.6 Results and. Discussion
In Table 8 the electrical conductivity values for various alkali and alkaline earth metal chloride solutions in HPO2F2 are given and they are plotted against the concentration expres• sed as a molality in Fig. 12. Included in the figure are the electrical conductivity curves for the corresponding difluoro- phosphates (shown as a solid line) and as can be seen the conduc• tivity curves for the difluorophosphates at low concentrations are virtually identical with the experimental points of the chlorides. However, at high concentrations some deviation Is TABLE 8
Specific Conductances of Some Inorganic Electrolytes in Difluorophosphoric Acid at 25°
KC1 NaCl CaCl2
102m. lO^X 102m. lO^ 102m. lO^X ohm~lcm7-'- ohm-lcm.-i ohm**lcm.
0.000 2.780 0.000 2.789 0.000 2.783 2.870 3.943 3-550 5.200 1.637 3-924 6.711 6.117 12.30 11.22 3.692 5.510 12.72 9.397 24-08 18.00 6.874 7.935 21.80 14.11 33.30 25.50 11.96 10.75 34.26 21.18 55.78 33.55 17.24 13.21 51.39 29.80 22.68 15.40 67.52 38.17 30.72 17-90
tetramethyl-
KNO3 ammonium chloride BaCl2 0.000 2.616 0.000 2.652 0.000 2.757 0.8532 3.147 3.061 4.513 2.661 5.450 3.023 5.391 7.357 8.237 6.935 10.86 6.052 8.142 11.44 12.57 10.93 15.52 10.95 12.35 16.93 19.59 16.81 22.20 15.85 16.95 21.28 26.11 21.04 21.31 26.28 34.67 27.44 26.00 30.72 43.06 34.69 50.86
KgCO^ tetraphenyl- AsF arsonium chloride 3 not soluble 0.000 2.622 0.000 2.515 0.7154 3-152 0.3773 3.093 1.577 4-499 1.231 4.744 2.835 6.368 2.198 6.946 •5.280 9.915 3.731 11.02 9.220 16.50 4.963 14.19 15.00 23.75 6.528 18.29 20.60 32.78 8.039 22.19 25.00 • 41.18 9.344 24.70
57
observed. The most reasonable explanation of these results is
that complete solvolysis of the chloride occurs to give the metal
difluorophosphate and HC1. Over a wide concentration range the
HC1 has no effect on the.conductivities of the difluorophos•
phates (i.e. HC1 is a non-electrolyte), however, at high con•
centrations the HC1 could cause a change in the properties of
the solution such as viscosity, density and dielectric constant
which could increase the mobilities of the ions present or in•
crease the degree of dissociation of ion pairs and thus produce
a higher conductivity.
It is well known that in solvents of low dielectric con•
stant where ion association is important, salts of large cations
show the greatest degree of dissociation19. To further test the
assumption that metal difluorophosphates are strongly associated
in HP02F2, conductivity studies on (CH^)^NCl and (C^H^)^AsCl were made. Since ionic chlorides apparently undergo complete
solvolysis (see above) to the corresponding difluorophosphate
and hydrogen chloride these complex salts will react with HP02F2
according to the equation,
R^MCl + HP02F2 —* R MP02F2 + HC1 (3.3) '
and it' is thus important to know 'what effect the HC1 produced will have on the conductivity of the solutions. The studies on
the alkali and alkaline earth metal chlorides have shown that
at low concentrations the effect of the HC1 is small and the con•
ductivity observed is essentially that due to the corresponding
difluorophosphate. It may be assumed that the conductivity 58 curves for the R^MCl compounds are essentially those of the cor•
responding difluorophosphates. While (CH3)^NP02F2 exhibits con• ductivity values about the same as those of metal difluorophos• phates the compound (C^H/^^AsPO^Fg is much more highly conduct• ing. It is unreasonable to assume that these high conductivities are due to a greater mobility of the (C^H^)^As+ cation compared to the metal cations and must be due to greater dissociation of this salt, confirming the former conclusion that the metal di- fluorophosphates are weakly dissociated.
In Fig. 13 the conductivities of potassium nitrate and carbonate are shown relative to KC1 and KPO^Fg. The conductivity produced by KgCO^ solutions is more than that expected for the solvolysis of the carbonate to 2KPO2F2 according to,
KgCO^ + 2HP02F2 ~~* 2KP02F2 + H2C03 (3.4)
H2C03 —» H20 '+ C02
Considerable effervescence was observed during the run presumably due.to the evolution of carbon dioxide. The water that is thus produced should be protonated to E^0+ in this acidic solvent.
Thus the observed conductivity for the carbonate solution should be greater than twice that of a similar concentration KPOgFg solution. However, it is not known If water behaves as a simple
base in the HP02F2 solvent system; it is possible that equilibria of the type,
H20 + HP02F2 5=* H2P03F + HF
+ H P0 F + HF H30 + P02F2" ^ 2 3
could occur. Protonation of H2P03F or HF produced in such reactions- could account for the fact that the observed conductivity is 6£ 60 greater than twice that for KPOgFg. 19 Potassium nitrate is known to react with sulphuric and fluorosulphuric acids to produce the nitronium ion,
+ KN03 + HA ^ HNO^ + A~ + K (3.5)
+ HNO^ + 2HA ^ N02 + 2A" + H^O*
In difluorophosphoric acid the conductivities observed for KKO3 are much greater than those observed for similar concentrations of KPO2F2, thus indicating that some reaction in addition to the simple solvolysis reaction
+ KNO^ + HP02F2 —• K + P02F2~ + HNO^ (3.6) is taking place. It is not known if nitric acid is stable in
HP02F2, nor whether it would act as an acid or base. It is doubtful if protonation of nitric acid would occur to any great extent.
+ HNO3 + HP02F2 ^ H2N03 + PO^" (3.7) During the addition of nitrate brown fumes were observed and the solution was coloured brown, it thus appears that the nitric acid is unstable and decomposes according to,
+ 2HNO3 + HP02F2 ^ 2N02 + P02F2" + H30 + 02 (3.8) The observed conductivity values which are greater than those
exhibited by KP02F2 at the same concentrations could be explained by the presence of this reaction.
It is impossible, at the moment, to be more quantitative, concerning these studies as the ionic mobilities of the species postulated and the degrees of dissociation of the ion pairs are not known. It could be, for example, that the X versus n curve for potassium carbonate is just less than that of 3 x X
versus n. curve for KP02F2 due to the fact that the ionic mobility of the H^O is less than that of K or the degree of
+ dissociation of H30 P02F2- is less than the degree of dissocia•
+ tion of K P0oFo~. 62
CHAPTER IV
Acids and Acid-Base Reactions
4.1 Introduction
Few subjects in chemistry have excited more interest and more resultant controversy than the subjects of acids and bases'^. Acids and bases are regarded as mutual opposites which in general lose their defining properties when brought into, con• tact with each other. The solvent system concept is based upon the formulation of acidic and basic entities in a solvent as those species which result from the self-ionization of the solvent, e.g.
4 2H20 ^ H^O" + 0H-
In difluorophosphoric acid any solute which causes an increase in the dif luorophosphoric acidium ion, H^jPO^F^, ion concentra• tion will be considered an acid, and correspondingly any solute
which causes an increase in the difluorophosphate ion, P02F2~, concentration is a base. 22 It has been shown that in the sulphuric acid solvent
system HP02F2 is a base, HSO^F an acid and CF^COOH a non- electrolyte and it is likely, therefore, that in the difluorophos• phoric acid system, the protonic acids HSO3F, H2SO4 and CF3COOH will act as acids according to (1.5). The expected relative
order of acidities in HP02F2 is HSO^F > H2SO^ CF^COOH. Lewis acids of the type SbF^, BF^ and AsFc; may also prove to be acids by being anion acceptors according to equation (1.6). It
should be possible to carry out acid-base neutralization reactions in HPO2F2 as in any other amphoteric solvent. The neutraliza• tion reaction is simply the reverse of the solvent autoprotoly•
sis i.e.,
+ H2P02F2 + P02F2" — 2HP02F2 (4.1) -f _ In sulphuric acid the. autoprotolysis ions H-^SO^ and HSO^ have very much higher mobilities than any other ions and thus neutral ization reactions are conveniently followed by electrical conduc timetric methods4 . As there seems to be some uncertainty as to whether the autoprotolysis ions of HPO2F2 exhibit abnormal mobility, it is anticipated that there xvill be considerable difficulty in the detailed interpretation of conductivity titra• tion curves. Further, any ion-pairing will cause even more dif• ficulties in the detailed interpretation of the data.
4.2 Experimental ' • A) Preparation and purification of materials.
Fluorosulphuric acid: commercial HSO3F obtained from
Allied Chemical Co. was.double-distilled at I640 and atmospheric 10 pressure. Trifluoroacetic acid: Commercial CF^COOH from Eastman
Kodak Co. was double-distilled at 72° and atmospheric pressure. .• Antimony pentaf luoride: commercial SbF^ obtained from
Alfa Inorganics Inc. was double-distilled under vacuo.
Sulphuric acid: 100$ HgSO^ was made from 98$ "AnalaR" sulphuric acid and dilute oleum according to the method of minimum conducting acid due to Gillespie et al.*-4-' 6Z(.
B) Electrical conductivity.
The cell used is that shown in Fig. 2 and the procedure adopted for the conductivity runs has been discussed previously. As the solutes are liquids, the microburette was used to make the additions/ However, in the case of antimony pentafluoride a concentrated solution of SbF^/HPO^Fg was made and this was added by means of the microburette. The solution was made by distilling SbF^ under vacuo into a clean, dry tube which was then reweighed and transferred to the dry box. A suitable amount of freshly distilled HPO2F2 was added to the tube which was then reweighed. The resulting solution usually had a concentration of 3-5 molal. The tube was shaken vigorously for several minutes and left for 30-40 minutes to ensure the solution was homo• geneous. The mixing of the two. liouids 'was exothermic. During the shaking of solution A in Table 9 the tube was cooled in liquid nitrogen. A sample of the resulting solution was drown into the microburette and used for the conductivity run.
In all cases after addition of acid was complete, base was added by means of the injector (Fig. 3) pnd the conductivi• ties of the resulting solutions recorded. In the cases of HgSO^, HSO^F and SbF^ white precipitates were obtained on the addition of base. The licuid was removed by filtration and the precipi•
tate washed with HP02F2, dried by pumping off the excess acid at room temperature and stored over phosphoric oxide in a desic• cator. C) Nuclear magnetic resonance.
The 1H, 19F and 31P spectra of solutions of SbFc, in
HPOgFg were obtained on a Varian HA100 high-resolution spectro• meter operating under the conditions described previously.
Calibration of the spectra was obtained by locking on to a suit• able peak e.g. in the phosphorus spectrum the central peak of the triplet was used, and scanning over 1000 cycle sweep widths.
The solutions studied were of a higher concentration (usually
2.5 molal) than those studied conductimetrically and were ex• amined at room temperature and at around -70°.
4.3 Protonic acids: Results and discussion
In Table 9 "the results of the conductivity measure• ments on solutions of the protonic acids in HPOgFg at 25° are given. In Fig. 14 the specific conductances are plotted against the molal concentration, m. The conductivity curves of
HSO3F and H^SO^ show an initial flat portion followed by a linear increase in conductivity with concentration. The flat portion of the curves could be attributed to titration of basic impurity in the difluorophosphoric acid. However, it is possible that in dilute solutions the acids are quite weak but at higher concentrations where solute-solute interaction is possible the polymerized forms of the acids are much stronger and are capable of greater proton donation to solvent molecules.
At the same concentration the conductivities increase in the order CF^COOH< HgSO^< HSO3F. Defining acid strengths
H in terms of the concentration of the free ions HoP0oFo and A" TABLE 9 Specific Conductances of Some Acids at 25
CF3COOH H2S04 HSO3F 102m 102m 10^ 102m 10^K ohm -Lcm. ohm-^cmr-1- ohm-lcmrl 0.000 2.710 0.000 2.837 0.000 2.485 0.1514 2.729 0.2902 2.805 0.2830 2.525 0.5552 2.772 0.5223 2.B19 O.566O 2.540 1.393 2.860 0.8560 2.795 O.848O 2.578 2.645 2.999 1.407 . 2.770 1.4-15 2.698 4.502 3.201 2.162 2.773 1.697 2.732 7.107 3.486 2.916 2.811 2.420 2.885 9.994 3.810 3.671 2.369 2.850 2.960 15.01 4.368 4.837 2.982 20.02 4.915 6.239 3.156 0.000 2.553 24.02 5.347 8.183 3.444 0.1450 2.556 28.08 5.773 10.65 3.851 0.5517 2.537 32.11 6.191 13.09 4.280 1.277 2.682 36.15 6.600 16.31 4.874 2.338 2.879 40.19 7.001 19.28 5.433 3.938 3.222 Addition of KP02F2 23.65 6.336 6.166 3.785 Total 28.29 7.337 8.676 4.466 0.9944 41.18 7.355 34.18 8.679 12.32 5.591 3.294 43.48 8.353 40.10 10.07 5.990 46.18 9.713 49.33 12.45 0.000 2.580 59.11 15.13 11.80 5.020 Addition of KPOo2*F2 17.70 7.460 Tota 1 23.60 9.790 0.8763 59 .99 15.05 29.50 12.19 2.939 62 .10 14.40 42.00 18.86 7.455 66 .57 11.35 53.10 24.60 .46 69 IO .57 10.05 Addition of NH2P02F2 Total V 6.79 59.39 19.33 22.9 76.00 10.25 45.6 98.70 4.48 61.3 114.90 8.16 TABLE 9 (cont'd)
A B
SbF^ SbF5
Initial conc'n of Initial conc'n of
SbF5/HP02F2 = 3.733m. SbF^/HPO^ = 5.484m. 102m 10/lX 102m 10*11 ohm~lcmrl ohm~lcmr- 0.000 2.732 0.000 2.574 2.097 3.079 0.1654 2.595 5.378 4.476 0.3221 2.626 10.18 7.080 0.5361 2.674 16.10 10.96 O.8488 2.741 Added KPOoFo 1.201 2.832 Total 1.615 3.027 11.22 27.32 6.035 2.076 3.186 21.68 37.78 6.040 2.635 3.415 23.72 39.82 7.286 3.342 3.675 4.201 4.051 5.338 4.686 7.356 5.461 9.122 6.272 II.64 7.531 14.13 8.781 17.49 10.53 21.43 12.69 25.39 14.96 29.48 17.36 33.43 19.27
Added KP0?F9 Total 3.042 36.47 17.69 8.247 41.68 14.95 19.21 52.64 9.427 30.25 63.68 7.756 32.69 66.12 8.257 38.64 72.07 9.900 46.38 79.81 12.95
69 in solution, the higher conductivities exhibited by HSO3F solu• tions suggest this acid is the strongest acid of the three in difluorophosphoric acid solution. In sulphuric acid solutions 22 other workers have shown that HPO2F2 is a weak base of the
H2SO4 system and the results reported here shov/, as would be ex• pected, that in the HPO2F2 system. R^SO^ is only a weak acid causing little protonation and hence producing few ions and a low conductivity. Also from the sulphuric acid work it. has been suggested that CF^COOH is a weaker base than HPO2F2, in fact it is a non-electrolyte, and hence it seems likely that in the
HPO2F2 system CF^COOH'will be a weak acid, of comparable strength to R^SO^. From the conductivity curves in Fig. 14 it appears that CF3COOH is ionized to a smaller extent than the inorganic oxy-acids.
The effect of adding base to solutions of the protonic acids was followed conductimetrically and the results are sho\vn in Fig. 15. For H^SO^ and HSO^F solutions a white solid separ• ated and the conductivity decreased. In the case of HSO^F the
base,-NH4P02F2, was added until the conductivity increased again;
the base was added to a 0.531 molal solution of HS0^F/HP02F2 and a minimum was obtained at a total molality of 1.05. It therefore appears that the base neutralizes the HSO^F in a 1:1 reaction,
+ HSO3F + HP02F2 — H2P02F2 + S03F~ (4.2)
+ NH4P02F2 ^ P02F2- +NH4
The white solid \^as identified as MH^SO^F by comparing its infra• red spectrum-with that reported by D. W. A. Sharp.45 The observation
71 of the production of the salt NH^SO^F, of course, confirms the previous assumption that HSO-jF behaves as an acid in HPO2F2 solutions.
Addition of the base KP02F2 to the H2S0^/HP02F2 solution caused the precipitation of KHSO^ which was identified by its melting point (210 )° and its infrared spectrum (compared to that of an authentic sample of KHSO^ - Baker A. R. Grade).
+ HgSO^ + HP02F2 ^ H2P02F2 + KSO^" (4.3)
P0 F + K+ KP02F2 ^ 2 2~
Again, the observation of the production of KHSO^ confirms the
assumption that H2S0^ behaves as an acid .inlIP02F2. Addition of
the base KP02F2 to the CF^COOH/HPC^Fg solution causes the conduc• tivity to increase and no precipitate is formed. As no product
from the reaction of CF3COOH with KP02F2 in HP02F2 was identified it is not possible to say for certain that trifluoroacetic acid
is behaving as an acid in HP02F2. It is possible that CF3COOH
is.behaving as an acid and the salt formed on adding KP02F2,
CF3C00K is soluble in HP02F2« Unfortunately no studies on the solubility of this salt in HPO2F2 have been made as yet. The increased conductivity observed on addition of KPO2F2 to the solutions of CF^COOH could be due to a greater degree of dis• sociation of.CF^COOK over that of CF3COOH.
It had been hoped that the acid-base titrations would give some indication as to whether abnormal conduction is shown by
the autoprotolysis ions of HP02F2. In H2S0^ conductimetric acid- base titrations have confirmed that the H^SO^ ion has an abnor- 43 mally high mobility. Addition of KHSO, to a solution of an acid (e.g. HB(HSO ) ) causes a marked decrease in the conduc•
tivity as the highly mobile H^SO^4" ion is replaced by the relative
ly poorly conducting K""" ion. However, in HPO2F2 addition of base
causes precipitation in the cases of HSO3F and H2S04 and thus a
a loss of ions due to the insolubility of the salt. The decrease
in conductivity cannot be attributed to the replacement of the
+ + H2P02F2 ion by the M ion.
4.4 Solutions of antimony pentafluoride: Results and discussion.
From a study of the chemical reactions of a large number
of fluorides in solution in hydrogen fluoride, Clifford et al.4^
concluded that some, solutes exhibited acid behaviour by acting as fluoride ion acceptors. The pentafluorid.es of Sb, As and P were found to be acids which decreased in strength in the order
SbF^> AsF^>PFy The first quantitative study of solutions of
SbF 5 in HF was made by Kilpatrick and Lewis*"7 who studied the electrical conductivities of solutions of SbF^ and also of KF, 48 NaF and NaSbF^. Hyman et al. investigated the infrared and Raman spectra of SbF5 solutions in hydrogen fluoride and obtained evidence for the formation of the SbF^" ion. Their work also in• cluded a.study of the electrical conductivities of the HF-SbF^ system over the whole range of composition, but not of very 49 dilute regions. Hyman and Katz concluded that in HF the net result .of the addition of a ILuoride ion acceptor is precisely the same as the addition of a proton donor, namely an increase in the concentration of solvated protons,
+ SbF.5 + 2HF = H2F + SbFg" . (4.4) Gillespie and Moss'^ were able to confirm the conduc• tivity results of Hyman et al. at high concentrations, however, at low SbF^ concentrations the conductivities obtained were much higher than those of Kilpatrick and Lewis. They also studied the 19F and -4-1 n«.m.r. spectra of a number of SbF^/HF solutions.
In the case of the proton n.m.r. the solutions gave only one signal even at low temperatures, whereas the fluorine spectrum contained several peaks which were assigned to a dimeric ion
F - S"b2 ll"> SbF^ and HF. The Sb2F]_]_~ ion was believed to form at high SbF^ concentrations due to the reaction,
SbF6~ + SbFr Sb2F1:L- ' (4.5)
In view of the fact that SbF.^ is itself highly polymerized, it was also suggested that higher polymeric ions, such as Sb^F-^g"
and Sb^_F2-]_~, may be present in the more concentrated solutions. 51 52
Woolf and Barr both investigated solutions of anti• mony pentafluoride in fluorosulphuric acid and Woolf suggested two possible modes of ionization for this solute, namely
+ SbFr + 2HS0oF =^ SbFrSCkF" + H9SO,,F .. 1 5 3 (4.6)
+ SbFr •+ 2HS03F SbF6~ + SO^ + H2S03F 2
Woolf suggested that equation 2 is more probable than 1 because the structure SbF^SO^F" would require abnormal five-fold co• ordination of sulphur while the alternative structure SbF^SO^" would require seven-fold coordination of antimony. Barr, however, pointed out that if the fluorosulphate group is bonded to an• timony via oxygen, no abnormal coordination is required. 53 Further work by Thompson et al. involved conductimetric, cryoscopic and nuclear magnetic resonance studies on solutions of. S0F5, SDF4SO3F and SDF5 - SO3 mixtures in fluorosulphuric acid. They were able to shov/ that a series of acids exist with
sbF S0 F the general formula H ( 5_n( 3 ) 1+n) where n^O, 1, 2 and 3..
Dimeric and probably higher polymeric forms of these acids were suggested to be present in the solutions also and n.m.r. studies showed that polymerization occurs through fluorosulphate bridges.
From the results of the SbF^/HF and SbF^/HSO^F systems
it was expected that in the HP02F2 solvent system SbF^ would prove to be a strong acid and a difluorophosphate anion acceptor ac• cording to,
+ . SbF5 + 2HP02F2 ^ SbF5P02F2"'+ H2P02F2 (4.7)
As the concentration of the SbF^ increases, it is possible that more highly polymeric species may be formed as found in the HF and HSO^F solvent systems.
In Table 9 the results of the conductivity measurements
of solutions of SbF^ in HP02F2at 25° are given and plotted against molality, m, in Fig. 14. The SbF^ solutions are more
conducting at the same concentrations than HSO^F and H2S04 solutions which is in agreement vriLth the relative order for the
acid strengths, SbF^> HSO^F^gSO^. Addition of KP02F2 causes the conductivity to decrease and a white solid to separate. In
Fig. 16 two sets of data are shown and in the more dilute SbF^ solution (curve A) the addition of base produces a minimum in the conductivity at approximately the end point for a 1:1
reaction of SbFc and KP0oFo.
76
+ SbFr + HP02F2 - HSbF.P02F|'-!- HP02F2 ^ H2P02F2 + SbFrP02F2~ (4.8)
+ •KP0.2F2 ^ • P02F2- +./ K It is difficult to be certain exactly where the minimum in the
conductivity titration curve occurs; however, it appears that,
from the titration of the more concentrated solution of SbFr 5 (curve B), the minimum occurs at a mole ratio of base to acid of
slightly less than one. Also, the conductivity at the minimum is greater than that of the pure solvent. These observations lead
one to suspect that reactions-in addition to (4-3) are taking place and this conclusion is confirmed by the n.m.r. studies to be presented later.
The n.m.r. spectra of various SbFr_/HP02F2 solutions were taken at -70° and 30° for19 F, 31P and hi and these are listed below: spectra molality temperature : : 3.63 30°, -65°
19F ' 3.63 30°, -65° shown in Figs. 17 & 18 !9F 2.31 30°, -65° details shown in Figs.20,21,22 & 23
31P 2.31 30°, -70° details in Fig.25 3'lp 2.62 30° .shown in Fig. 24 31P 2.27 30°, -70°
In the n.m.r. spectrum only one proton si gnal is observed at
30° and -70° thus indicating that ':rapi d proton transfer must be occurring between the various species present in solution.^
Molecules such as H(SbF5P02F2) and its anion SbF5P02F2" will be in equilibrium according to,
--alternate form is SbF^HPOgFg). 77
HSbFrP02F2 + HP02F2 ^ SbF5P02F2~ + H2P02F2 (4.9)
because of the rapid proton transfer only a single -^P and19 F
n.m.r. spectrum will be observed for these two antimony species.
For simplicity in the remainder of this discussion, the spectra will be assigned to neutral entities only, it being understood
that the spectra are probably combined spectra of the anions and
their corresponding neutral protonated molecules.
Considering first the n.m.r. spectra, the general
features of the two solutions studied (3.63 and 2.31 molal SbF,-) were the same; both spectra were obtained on 1000 c/s sections and in the case of the more concentrated solution also on a
single scan which is shown in Fig. 17. Only the spectra of the
2.31 molal solution will be discussed in detail as in the spectrum
of the 3«63 molal solution several peaks went off scale. At room temperature'in the region of the spectrum due to fluorine bonded to phosphorus several sharp peaks were observed while in the region due to fluorine bonded to antimony a few very broad resonances were present. These latter resonances are presumably broadened by fluorine exchange at room temperature. At -65° some broadening (cf. Fig. 20) was observed in the peaks due to fluorine bonded to phosphorus (probably due to increased vis•
cosity of the solutions) while fine structure in the region of fluorine bonded to antimony was observed. In Figs. 17 and 18 the entire 19F spectra at 30° and -65° respectively are shown.
The large doublet (labelled D and I in Fig. 17) is assigned to
solvent and on dilution with HP02F2 to produce a SbF^/HP02F2
solution of about 1 molal concentration the peaks D and I 19 Fig. 17 F Spectrum of a 3-63 molal SbF_/HP0_Fo Solution at 30° N.M.R. 5' 2 2 D
Fig. 19 F N.M.R. Spectrum of a 7.0 molal SbF5/HP02F2 Solution at 30°
CO- o (inset Fig. 17) clearly shov/ a large increase thus confirming their assignment to solvent. In Figs. 20-22 details of the .two regions for the 2.31 molal solution are given. The 3.000 c/s side bands produced by the spectrometer unfortunately confuse the overall spectra; however, using 1000 c/s sweep widths and the
"lock-on" technique eliminates these side bands. Unfortunately it is not possible to correlate areas under the peaks or peak heights in Figs. 17 and 18 with the number of fluorine atoms in the proposed species assigned to a peak, as in the scan of the complete spectrum overlap of modulation side bands alters the peak heights and areas. On the scan of the 1000 c/s sweep widths spectrometer conditions were changed from one spectrum to another so that a peak could be magnified by several times and hence its height or area would not give a true representation of the number of fluorines. A final difficult}/- lay in the fact that the solu• tions were so dilute that in order to record any- small peaks the magnification had to be so large that the peaks due to solvent went off scale and could not be measured. In Table 10 all the peaks observed, their multiplicity and coupling constants (J) are given; the chemical shifts are measured in cycles per second from the centre of the doublet due to the solvent and are the average values for all the 19F spectra examined.
At high field strengths, in the region of the spectrum due to fluorine bonded to antimony, the spectrum is composed essentially of a doublet-quintet pattern (labelled M and P in
Figs. 18, 21 & 22a) and a triplet-triplet pattern (labelled TABLE 10 ^"°T and ^lp Chemical Shifts* and Coupling Constants for Some Complex Antimony-Fluorine Species
Species 19F Chemical Shifts*
SbF5P02F2 (I) Doublet(CH) J Doublet(M) Quintet(P) AS ,J -50 933^ 2708 4254 1546 103
53 SbFrS0oF Doublet Quintet AS J 5 3 8450 100 10123 1678
(SbFjp (P02F2) ) • Ila Doublet(AF) J Triplet(K) Triplet(N) J x -98 994 • 1735 3190 114 lib Doublet(BG) J Singlet(L) ASKN A&K L -87 990 1994 1445 259 53 (SbF4(S03F))2^ Ha Triplet(F) Triplet(P) J 7832 9275 126
lib Singlet(H) ASFP A$FH 8129 1443 297 Unknown Singlet(0) 3510
POF3 Doublet(EJ) J 469 1060
* Measured from the centre of the doublet due to solvent
CO- TABLE 10 (cont'd)
•19F and -^P Chemical Shifts* and Coupling Constants for Some Complex Antimony-Fluorine Species
Species 31p Chemical Shifts*
SbF,P0oFo . Triplet(CHL) J ? * * • 273 986
o (SbF.(P0 R)) lla Triplet (DI,M1 ) J 4 * A. x 320 . 988
lib Triplet(EI9N) J 327 989
•POF, Quartet(AGKO) J ^ 583 1059
* Measured from central, peak of solvent triplet
co. 84
19 Fig. 20 F N.M.R. Spectra of the P-F Region for a 2.31 molal
SbF5/HP02F2
Peak E (Fig. IS) is used as
' a 'lock"
30° • -70'
A B
J
460 471 516 548 • 514 -500 -465 -410
H c/s-*
Fig. 22(b) i9F N.M.R. Spectrum in the Sb-F .Region for a 2.31 molal SbF^/HPO^ Solution
at -70° (peaks N & 0)
i
Peak E (Fig. 18) is used as a "lock" K and N in Figs. 18, 22a & 22b) and two single peaks (labelled. L & 0 in Fig. 18). The doublet-quintet pattern indicates that the expected reaction,
SbFr + HP02F2 -> HSbFrP02F2 (4.7) does indeed take place. Assigning structure I to the species
HSbFrP0„F„
I
the peaks (M) are due to the four equivalent fluorine atoms,
F]_, split into a doublet by spin-spin- coupling with F2. The
quintet (P) is due to F2 split'by the four equivalent fluorines,
F-^. The doublet and quintet are shown in more detail in Figs.
21 and 22 respectively. The observed coupling constants and
chemical shifts compare remarkably well with those observed for the species HSbF^SO^F^. The peaks (C & H) in the fluorine on phosphorus region (Figs. IS & 20) are tentatively assigned to the fluorine atoms F^.
The remainder of the peaks observed in the ^9F spectrum must be due to species produced in reactions other than (4.7).
The peaks labelled E and J (Fig. 18) are assigned to POF^. This assignment is based on the observed coupling constant of 1060c/s
(c.f. literature value of 1055c/s)"^+ and the fact that a quartet exhibiting the same coupling constant appears in the -^P spectrum 89
(see later). The two triplets (K and N) and one singlet, L, are
assigned to polymeric (SbF^POgFg).^. Cis P02F2 bridging in this
molecule (structure lla)
F6 F6 F6 F6 •\ / \ / \ F4 F4 ^p\. F4 ^ °\ 1 0. I °\ I 0 SbCT _^ J^Sb^f
F5 I ^F5 F5 | ^F5 F5 | ^ F5 F F 4 4 F4
lla
would give rise to the A2X2 spectrum observed in the fluorine on 55
antimony region/'' while trans P02F2 bridging (lib) will give rise to the singlet, L. Fg. ^-Fg F 7 ^ PC^ F7 F 0 7 ^ I ^ 0 I ' F7 ^TSb J^sh^T ^ 0^ ^0 I F7 F7* I 0. ^-0 ^:PCT f7 1 f7 J: Fg ^Fg F^ ^Fg
lib
The two triplets, K and W, (N is shown in more detail in Fig. 22)
are assigned to the fluorines F4 and Fr, and the singlet, L, to Fy.
Peaks A & F and B & G in the fluorine on phosphorus region are
tentatively assigned to FQ and Fg, respectively. The assignment of-these peaks is based largely on analogy of the observed spec• trum with the compound (SbF^SO^F^ which has both cis and trans, bridging fluorosulphate groups. The chemical shifts and coupling constants'^ observed for (SbF^SO^F)^ and (SbF^PC^F^^ are compared in Table 10. 90
A possible explanation for the formation of POF^ and
(SbF^P02F2)x in these SbF^/HP02F2 solutions is that at high
concentrations of SbFtj the following reaction takes place,
sbF 2SbF5.HP02F2 SF* ( 4P02F2)2 + 2HF (4.10)
Similar reactions can be written for the formation of polymers
larger than the dimer. The HF may then undergo reaction with
HP02F2 to produce POF^,
HF + HP02F2 ^ POF3 + H20 (4.H)
The single peak, 0, in the fluorine on antimony region may be due to HF or possibly HSbF^ produced by the reaction,
HF + SbF5 -v HSbF/5 (4.12)
In Fig. 19 the 1°T n.m.r. spectrum at room temperature of
a more concentrated solution of SbF^ in HP02F2 (7 molal) is shown.
It Is apparent that the spectrum is quite complex with many dif- . ferent types of fluorine bonded to phosphorus In the P-F region
of the spectrum. At high SbF^ concentrations in HP02F2 it is
expected that polymerization of SbFrj, which occurs in the pure 56
liquid , will also occur in HP02F2 and thus species of a complex
nature will form. It is possible that species of the type
(SbF3(P02F2)2)x may also be produced and would thus complicate the n.m.r. spectrum. It is not known if these polymeric forms
will be further solvolyzed to species of the type HSbF^(P02F2)2
and HSbF3(P02F2)3. These reactions could account for the seemingly large amount of POF^ In the 19F spectrum in Fig. 19. At the moment an endeavour to interpret spectra of this type will not be attempted until the isolation and definite identification of some of these SbF^/PO^F^ species has been accomplished. It
should be mentioned that the conductivity results (curve B in
Fig. 16) of SbF^/KPO^Fg titrations support the ideas suggested
previously, as the end point for the KPO2F2 titration in the more
concentrated SbF^ solution is less than that expected for a 1:1 reaction it thus appears that not all the SbF5 is in the form of
HSbF5P02F2.
Finally, spectra were also run under high-resolution on
100 c/s sweep widths and in Fig. 23 the ~9F n.m.r. spectrum of
peaks F, G and H observed in the region due to fluorine bonded to
phosphorus are shown. Previously, the peaks C and H were assigned
to the fluorine atoms, Fg, In structure I and if these two equi•
valent fluorines were to couple with the four equivalent fluorines,
FJL, on the antimony atom then under high resolution a quintet
pattern should be observed. As can be seen from Fig. 23 peak H
appears as a quintet with a coupling constant, J, of 1.5c/s.
Peaks B and G were assigned to Fg in structure lib and under high resolution the coupling of Fg with the fluorines on antimony (Fy)
should give a nonet. Peak G is resolved into a triplet which is ,
probably the central lines of greatest intensity in the nonet;
the remaining peaks will not be observed on the scale used. Peaks
A and F were assigned to F^ in structure Ila, which will couple with the two sets of equivalent fluorines on antimony (F^ and Fc_) . and produce two quintets. In Fig. 23 it can be seen that these two quintets overlap and are not resolved.
The -^P•spectrum (Fig. 24) verifies the previous assign• ments and in Table 10 all the peaks observed, their multiplicity Fig. 23 yY High^Resolution N.M.R. Spectrum of the P-F Region for
2.31 molal SbF /HP02F2Solution at 30°
Peak DC (Fig. 18) is used as a."lock"
H
891.1 902.2 948 Fig. 24 3lp N.M.R. Spectrum of a 2.62 molal 30° F SbFt./HP02F2 Solution at B
hy-2 K M, W
yvyvl A
u vO Vo Fig. 25 Details of the 31p N.M.R. Spectrum of a 2.62 molal SbF^/HPO^ • Solution at 30° Peak F (Fig. 24) is used as a "lock"
K L
1095 1249 1301 1315 Hc/s
0 and coupling constants, J, are recorded, measurements are from the
centre peak of the triplet and are the average values of all the
spectra. As the -^-p n.m.r. spectrum obtained at +30° and -70°
•showed the same general features only the spectrum observed at
+30° is given (Fig.24) and details of a part of the spectrum aie
shown in Fig. 25. 54 The triplet BFJ is assigned to the solvent and the quar•
tet AGKO is assigned to POF^ (J=1059c/s). Three additional trip-'
lets are observed CHL, BI-L M and EI2N (details of K, L, M & N
are shown in Fig. 25) which may be assigned to P atoms bonded to
F^, and Fg fluorine atoms in structures I, Ila and lib respec•
tively. Under high resolution no fine structure was observed as ,
the peaks were broad and the background noise was of a high level.
4.5 Studies on KSbFcP0oFo
The white solid which precipitated from solution on the
addition of the base, KPO^F^, to a 0.16m solution of SbF^ in
HPOgFg was washed with HPC^Fg, dried by pumping off the acid and
stored over phosphoric oxide. No attempt was made to recrystal-
lize the solid and it was hoped that by washing the solid with
HPC^Fg the product would be essentially pure. The analytical data
obtained from A. Bernhardt ate given below:
obtained calculated for KSbFrP02F2 %F 36.98 37.25 %?• 8.51 8.69
From the fluorine and phosphorus analyses it appears that the
formulation,'KSbF^POgF2 is a reasonable one and together with the
conductivity and n.m.r. results it would seem that the species 96
SbF^P02F2 does in fact exist. An infrared spectrum of the solid
was taken in the manner described previously4, using KBr plates,
on a Perkin-Elmer 421 spectrometer and is compared to KPO2F2,
KSbF6, SbFrOH- and (CH3)2Sn(SbF6)2 below.
TABLE 11
Infrared Spectra of Various Inorganic Fluorine Compounds
57 KP02F2 KSbF6 SbFcOH" (CH3)2Sn(SbF6)2^ KSbF5P02F2
1332s
P0 1317s vasym PO vaSym 1310s 1210 w
Vsym PO 1148s H°0S ^sym PO
1020w
990m Sb-F 953m Sb-F
910m Sb-F 904m Sb-F
'asym 850s 840m VSym & PF & PF 83 2< 828m sym ^asym
7*2w
76ovw
73 7W
663 Sb-F 630 Sb-F 660s Sb-F 667s Sb-F
560 540 m
SsymPF2 503s 506m £svm PF2 482 495s m s = strong, m = medium, w = weak.
It appears that in the KSbFcP02F2 spectrum (col. 5) that the 97
infrared spectrum of the POgFg grouP (col. 1) can be discerned, also the band at 667 cra?^ can be assigned to Sb-F stretch-(cols. 58
2, 3, and 4). In (CH^)2Sn(SbF^)2 Goel has suggested that the medium intensity bands at 990 and 910 cm7^ are vibrational modes, infrared inactive for the octahedral SbF^ ion, which have become infrared active as a result of distortion of the SbF^~ group by the
(CH3)2Sn group. Of the remaining bands (1020w,• 953m, 904m and
737w) in the KSbFrP02F2 spectrum the medium peaks at 953 and 904 cm7 are in the same region as the bands observed by Goel for the dis• torted SbF^ and may be tentatively assigned to Sb-F vibrational
modes in the SbF^ part of SbF^P02F2 . An x-ray powder photograph of the solid was taken as described previously^" and visual com• parisons showed the obtained photograph to be quite different from that of KP02F2^ and KSbF^. It thus appears that the solid product is neither KPOgFg nor KSbF^ nor a mixture of the two.
To establish further the species KSbF^PO?^ a sample was decomposed under vacuo at 250° in a fluorinated nickel can which was connected to the simple vacuum line shown in Fig. 26. This consisted of two traps followed by connections for a molecular weight bulb, an infrared gas cell, and a manometer; further traps were included before the rotary pump. The gases evolved during the heating of the can were condensed out in the first trap
(-195°) which was then isolated from the can. The liquid nitrogen bath was removed from the first trap and placed around the second which was then connected by- opening the stopcock to the first trap. The contents of the first trap were allowed to warm to Fig. 26 Vacuum Line used for the Decomposition Cr ft of KSbFcP0oF~
i
M = molecular weight bulb
I = infrared gas cell N = nickel can 99
room temperature and were condensed out in the liquid nitrogen
trap which was then disconnected from the first trap. The cooling
bath was removed and the contents of the second trap were allowed
to warm up and expand into the molecular weight bulb, infrared-
gas cell and the manometer. The procedure was repeated.
The white solid remaining in the can after heating was
identified as KSbF^ by infrared (a single broad band at 665 cm!""1"
9 in the region 4000-500 cm7\ c.f. literature value'' of 660 crrC"*") and x-ray powder photograph (identical to that of an authentic
sample of KSbF^, supplied by Alfa Inorganics, Inc.). The results
are consistent with the decomposition of KSbFrP02F2 according to,
KSbF5P02F2 , —> KSbF6 + P02F (4.13)
The infrared of the gaseous sample was taken in a cell with Csl
plates and a path length of 90 nuns.. The pressure of the gas in
. the molecular weight bulb and the infrared gas cell was of the
order of 50 mms. of mercury. The molecular weight of the gaseous
product(s) was found by two determinations to be 98 -5. The
infrared spectrum taken on the Perkin Elmer 421 spectrometer is
given in Table 12 together with that of the reported spectra of
POF^^ and HPO2F2."*"2 It would appear that the gas obtained in « the first run is POF^ (molecular weight, 104) with possibly some a
SiF^ impurity. .Stafford remarks that in his study of KP02F2 at -1 1020 cm. the region is obscured by the SiF^ stretch, thus the
strong band observed at 1029 cm."""'" may be due to this impurity.
On comparison of the second run with the combined spectra
of POF^ and HP02F2 excellent agreement is obtained. The most
noticeable exception is the band also obtained in both experiments. 100
TABLE 12
Infrared Spectra of Gaseous Productts) From Decomposition of KSbF^P02F2 .
HPO F '2 2 POF3 - Gaseous Productts) P0F, + HP09F Vapour at 0° C 1st Run 2nd Run
2500- •3100 mb 2500-3150 mb 2500- -3IOO mb
1415 ms . 1416 s 1416 s 1416 ms
1332 vs 1332 s 1332 vs
1193 m 1191 m 1193 m
1081 vs 1079 vs 1081 vs
1031 vvs /1036 m.sh 1030 vs 1031 vvs V1029 s 990 "vs 983 s 990 vs 993 s 986 vs 983 s
881 ms 880 ms 881 ms
873 ms •872 w 872 s 873
820 w 825 w 820 VI
732 s 728 ms
535 m 528 m 535 m
501 m 501 m
484 m 485 ms 490 msh 484 m
473 s 478 s 477 vs 473 s
467
456 at about 730 cm, 1 which has no corresponding absorption in the 6l POF3-HPO2F2 composite spectrum. Corbridge and Lowe assign bands in this'region to P-F stretch, and report peaks at 721 cm."1,
728 cm."1 and 745 cm.'"1 for various monofluorophosphates. It seems likely, therefore, that some P-F species is present whos.e
remaining spectrum is hidden by the POF^- - HP0?F2 bands.
It thus appears that the complex decomposes to the hexa- fluoroantimonate and gaseous phosphorus oxy-fluorides, the major one being POF3. It is suggested that the POF3 is produced by
further decomposition of P02F which arises according to (4.14).
3P02F -» P0F3 + P205. (4.14)
No direct evidence for P2P5 was obtained, however, it Is likely that it would condense out in the vacuum system (sublimes at
300°)^ and give little indication of its presence in the infra• red spectra of the gaseous products and of the solid left in the can.
In the second experiment the vacuum system was.clearly
not properly dried and the P0F3 was hydrolyzed to give the ob•
served FiP02F2, according to,
P0F3 + H20 —> HP02F2 + HF (4.15) Hydrolysis of POgF may also cause the production of monofluoro•
phosphoric acid, •H2P03'F, unfortunately a. literature search yields no information on its infrared spectrum.
P02F + H20 —> H2P03F' (4-16)
Another possible product of the decomposition of KSbFr-P0oFp is P2O3FJ1 , however no evidence was obtained for its presence in the gaseous products. CHAPTER V
Nitrobenzene Solutions 5.1 Introduction
Gillespie et al. showed that in sulphuric acid the fol•
lowing order of acid strengths prevailed, HSO^F> H2S0^CF^COOH > HPO^Fg and the results in Chapter IV of the electrical conduc• tivity studies of various acids in difluorophosphoric acid con• firm this-order. Acidic behaviour has been mainly studied in 1 19
protonic media ' e.g. the solvents H20, HF, H2S0^, HC1, NH^ and CH^COOH; then to investigate further the factors influencing the
relative acid strengths of HP02F2, H2S0^ and HSO^F dilute solu• tions of these acids were studied in a non-protonic medium. It was decided to use nitrobenzene as the solvent as it is readily available and has itself been studied conductimetrically in
10 19 HSO^F, H2S0^ and HP02F2 (this work) and is a suitable solvent " for cryoscopic and electrical conductivity experiments. It was found that in HSO^F"^ as solvent nitrobenzene is a strong base
and is fully protonated whereas in HgSO^2 and in HF^9 it is a weak base. The reaction between nitrobenzene and the inorganic acids which provide the bulk solvent may be pictured as involving
initial formation of the adduct C^H^N02.HA followed by dissocia• tion into ions,
+ . . _ C6H5N02 + HA (C6H5N02.HA) C6HrN02H + A~ (5.1) A number of factors may influence the position of this equilibrium
it has been proposed that the reason for the differences In the
dissociation constants for these postulated C^H^NO^.HA adducts is 63
that fluorosulphuric acid is a stronger acid. , or better proton
donor', than either sulphuric acid or hydrogen fluoride. Propertie of the bulk solvent such as its dielectric constant and its
ability to solvate the Ions produced are also important factors.
Various workers have used nitrobenzene as a solvent
because of its convenient freezing point for cryoscopic studies;
Whittla^4 has investigated selenium tetrafluoride adducts in
nitrobenzene by electrical condtictivity and cryoscopy. Taylor" 65 and Kraus studied the electrical conductivity of various picrates and further work involved a study of substances which 66 fell into three main classes: (i) strong electrolytes,
(ii) weak electrolytes and (iii) weak electrolytes in which dis•
sociation into free acid and base takes place to a measurable
extent. However, it appears that little has been done to study
the effects of strong inorganic oxy-acids' in. nitrobenzene by
conductance and cryoscopy. One notable piece of work is that of
Oddo and Anelli who investigated the depression of the freezin
point caused by sulphuric and nitric acid, in a variety of sol• vents and they found that in nitrobenzene sulphuric acid existed mainly in the form of dimers whereas nitric acid remained as the monomer.
5.2 Experimental
A) Cryoscopy
The cryostat shown in Fig. 27 is similar to that used by
Whittla""1-. A single-hole cork was inserted into the B24.ground
glass socket and through this.the platinum resistance thermo•
meter was lowered. . The thermometer was held at the desired
position by pouring molten wax on the top of the cork and allow! 104 105 it to cool, it was held upright by a semi-circular wire holder which fitted into the cryostat just below the ground glass socket.
All ground glass joints on the cryostat were greased, with a narrow ring of silicone grease (Dow-Corning). Temperatures were measured by a Leeds-Northrup capsule platinum resistance thermo• meter which was calibrated by the National Research Council,
Ottawa. The resistance of the thermometer was determined on a
Guideline potentiometer (Leeds and Northrup, Type 43©3A) with a standard 10 ohm resistance (Leeds-Northrup).
The cryostat containing the teflon-coated, bar magnet stirrer was connected to a dry-air line by means of the B19 cone and flushed out with dry air for several hours. The cryo• stat was then transferred to the dry box and charged with a weighed amount of nitrobenzene (usually of the order of 60g.). The cryo• stat was removed from the dry box and cooled in a cold water bath
(5°). At' temperatures below the freezing point the rotation of the magnetic stirrer frequently caused the solvent to 'seed-out', thus it was necessary to discontinue stirring once the contents had cooled to the freezing point. The amount of supercooling xvas not allowed to exceed 0.5-0.9° (if a larger amount of supercooling were allowed, the quantity of solid which separates would be so large that the concentration of the solution would be greatly increased and give erroneous values for the freezing point de- 68 pression ). The cooling bath was removed and the air jacket
(J) was then placed around the lower part of the cryostat (C). The whole assembly was lowered into a second ice/water bath, which was kept at 2-4° below the freezing point of the nitrobenzene 106 solution. The magnetic stirrer v/as started and if freezing did not occur the solution was seeded by adding a small piece of platinum previously cooled in liquid nitrogen.
Readings were taken on the potentiometer, using a moving coil galvanometer as a null detector for fifteen minutes after the solution was seeded out. As soon as nucleation takes place, the latent heat of crystallization tends to raise the tempera• ture. The temperature cannot rise above the freezing point, but - if the cooling is too rapid it may. well never reach the true melting point. The freezing point for the solution was calculated from the maximum value observed with the platinum resistance ther• mometer. This value was attained three to four minutes after the solution was seeded out.
The liquid was allowed to warm up and the above procedure repeated. When reproducible values for the freezing point of the pure solvent had been obtained the first addition of solute was made. During the remainder of the experiment the solution was not allowed to rise above 8°. The freezing point of the solu• tion was determined in the same manner as described for the pure solvent; the amount of super-cooling was noted in each case.
For the additions of sulphuric and fluorosulphuric acids weighed amounts of nitrobenzene and each acid were mixed in weight-droppers in the dry box. Great care was taken to ensure a homogeneous solution. The dropper was weighed by suspending it from a single pan balance and an addition of the acid/nitrobenzene mixture was quickly made to the cryostat through the B19 cone; 107
the dropper was then reweighed. The freezing point of the result•
ing solution was found in the manner described previously. For
every compound at least three additions were made for each run
and for each addition the depression of the freezing point was
checked at least twice.
A correction was made to the solution concentration for the amount of solvent frozen out when the solution was seeded.
The weight of nitrobenzene frozen out Wn, is given by the ex- . 64 pression,
Wn = Ws Cp AT- AHf where Ws is the weight of solvent in the cryostat, Cp the heat capacit}' of nitrobenzene (1.52 gr^deg."! at 5°) , AT the amount of supercooling and AH^ the latent heat of fusion of nitrobenzene
(94.25 cal. g," ). 7 The molality of the solution, m, in equili• brium with the solid nitrobenzene was determined by the following equation,
• m = . 1000 M Tv^r where W is the weight of solute and M its molecular weight. No correction was made for the heat capacity of the cryostat.
The solutions of the acids in nitrobenzene were found to be unstable and darkened slc\-/ly; on standing for 24 hours the solutions turned black. The total time taken for each experiment was of the order of four to six hours and the decomposition during this time was not noticeable. Whittla0^ found that during cryoscopic experiments lasting three to four hours on selenium tetrafluoride adducts in nitrobenzene the decomposition was small enough to be neglected.
B) Electrical Conductivity
The conductivity measurements were made in nitrobenzene using the conductivity cell shown in Fig. 2. which was also used for conductance studies in HPO2F2. The procedure followed was similar to that adopted for difluorophosphoric acid. The weighed cell was connected to the distillation apparatus (Fig. 1) at K.
Purified nitrobenzene was added to the first distillation flask by means of a weight dropper. The apparatus was evacuated and the nitrobenzene was distilled under vacuo at a temperature of
40° into flask Y and finally into the cell. The pump was switched off and dry air was allowed to enter the system. The cell was quickly capped, weighed and placed on a stand in the oil bath set at 25°.
The acids were added to the cell in two ways; the first way involved the use of the microburette (Fig. 4). The pure solutes and microburette were transferred to the dry box, where the burette was rinsed out and then filled with the solute to be studied. It was then removed from the dry box and placed on the cell by means of the Teflon B19 socket adaptor. Additions were made by screwing up the plunger and recording the volume additions. The second way involved mixing known weights of the solute and nitrobenzene in the dry box. The weight dropper was filled with the resulting solution and reweighed. An addition was made to the cell and the dropper reweighed.
In both methods the solutions were mixed thoroughly after 109
each addition in the cell and the conductivity recorded; the
solutions were then mixed again and the conductivity redetermined until no further change was observed. In this way variations in
conductivity due to improper mixing were eliminated. As a control run the conductivities produced by solutions of tetra-n-butyl- ammonium bromide in nitrobenzene were determined (Table 14).
The results agree with literature values to within 2%.
C) Preparation of materials
AnalaR nitrobenzene was recrystallized from itself twice and then stored over molecular sieves, type 5^. It was then doubly distilled twice under vacuo in an apparatus similar to that used for the purification of difluorophosphoric acid. The nitro• benzene was collected in a flask which was quickly capped and transferred to the dry box.
Commercial benzil was recrystallized from alcohol'and was found to have a melting point of 94.7.-95.0°(95 ) .
5.3 Results and Discussion The freezing point of a liquid is lowered by the addition of a solute, and this loitering is directly proportional to the concentration of the solution for dilute solutions.
AT = KfiA where AT Is the freezing point depression, m the molality and
Kf the molal freezing point constant. In cases where the solute dissociates to form more than one particle a more general ex• pression can be used, 110
AT = Kf V m
where V is equal to the number of particles formed from each
molecule of solute. The values obtained from the depression of
the freezing point are given in Table 13 and in Fig. 2$ AT,
the freezing point depression, is plotted against concentration,
m, expressed in moles per 1000 g. of nitrobenzene. It is ap•
parent that the molal freezing point constant, Kf, is equal to
the slope of the benzil curve, assuming V" = 1. The value for Kf was found to be 7.00 per 1000 g. which agrees well with repor•
ted values by various workers (6.89-7.10°).^'"^
A) Fluorosulphuric acid solutions
It can be seen from Fig. 28 that HSO^F also appears to
cause a depression of the freezing point very near to that ex•
pected for V-l corresponding to one mole of particles per mole
of acid in the dilute region and slightly less, than one mole at the higher concentrations. It was expected that HSO^F which has been described as the strongest simple (i.e. isolable) acid
known ^ would extensively protonate the nitrobenzene and thus
cause twice the depression observed, because of the reaction,
r HSO3F + C6H5N02 ~* C5HrN02H' + SO3F" (5.2)
Moreover, the electrical conductivity results given in Table 14 and plotted against the molality of the solutions in Fig. 29 also give no evidence for this extensive protonation. Unfortunately, the reproducibility of the conductivity results v/as not good, reflecting the fact that the conductivity produced, by the solute is so low that traces of Impurity which may have been present in TABLE 13
Gryoscopic Measurements in Nitrobenzene
Fluorosulphuric Acid
T° AT0-• 102m T° AT0 102m 5.697^.005 , 0.000 5.707^.005 0.000 5.201-.010 0.4961.015 7.172 5.151-.010 0.556-.015 8.177 4.4797.045 I.2187.O5O 18.21 4.543x-010 1.164~-.015 16.88 3.120"-.020 2.577x.025 39.02 4.062±.020 1.645T-025 24.38 3.544i.020 2.1637-025 31.89 2.477-.005 3.230-.010 48.8O
Sulphuric Acid
T° AT0 102m T° AT0 102m
5.73 6-.-015 0.000 5.73 6i. 005 0.000 5.603±.010 0.133"-.025 2.655 5.500"i.005 0.236±.010 4-428 5.309-.010 0.427-.025 8.902 4.647-.010 0.889-.015 18.35 5.098^.015 0.638*.030 14.90 3.503^.040 2.233-.050 53.40 4.52 ±.011 1.154"-.026 24.55
Benzil
T ° AT° 102m T° AT° 102m
5.692±.020 0.000 5.707±.005 0.000 5.633±.020 0.059-.040 0.9978 5.657J-".005 0.050±.010 I.O84 5.564^.005 0.1281". 025 2.054 5.324±.014 0.3 83 ±.019 5.675 5.397±.005 0.295±.025 4.904 4.985±.000 0.722±.005 10.53 4.430±.005 1.277±.010 18.13
T is the average value of the freezing point
AT0 is the average value of the freezing point depression caused by the addition of solute
m is the concentration of the solute expressed in moles per 1000 g. of nitrobenzene Fig. 28 Depression of Freezing Point (AT) for Various Solutes in Nitrobenzene
0 ' ' 10 20 .30 102 x molality, m the solute, or produced by minor side reactions between solute and
solvent, or may have entered the solution during handling, affect the results to a considerable degree. Nevertheless it is possible to estimate the equivalent conductivities produced by solutions of HSO^F and these are presented in Table 15. The uncertainty in the values, given in brackets, was determined from the experi• mental scatter. .In spite of the approximate nature of these equivalent conductivity values it is clear that they are lower even than those obtained for weak electrolytes such as trimethyl- hydroxyammonium picrate which has a dissociation constant of
1.7 x 10~^ mole 17"^ in nitrobenzene.0^
From the results it may be concluded that HSO^F dissolves in nitrobenzene essentially as a non-electrolyte. If any dis• sociation occurs the dissociation constant is probably less than r i 10
2 x 10 J moles 1. . When HSO^F provides the bulk solvent , however, the dissociation is so extensive that the equilibrium constant cannot be measured. Since dissociation constants as 2 -1 high as 1.3 x 10 moles 1. have been measured in HSO^F it is reasonable to assume that the dissociation constant of nitro- 2 benzene In this acid solvent is greater than 2 x 10 . Hence, 7 the dissociation constant decreases by a factor of 10' in going from solutions in fluorosulphuric acid to solutions in nitrobenzene.
Nitrobenzene has a dielectric constant of 34.5 and that of HSO.-jF has been estimated to be 120,"^ thus the difference in the dielectric constants will account to some extent for the dif• ference in the dissociation constants observed. The effect may 114
TABLE 12,
Specific Conductances of Some Electrolyte Solutions in Nitrobenzene at 25°
Tetra-n-butylammonium Fluorosulphuric acid (microburette) bromide
2 2 7 io :m lO^ 10 m 10 X ohrfl-lcmr-'- ohm-lcmrl 0.000 4.607 0 .000 6.979 0.6194 1766 4 .580 • 29.60 0.8927 2354 9 .119 63.59 1.214 2969 13 .61 107.2 1.513 3533 22 .49. 206.9 31 .21 332.2 0.0000 2.644 39 .76 476.8 0.04980 182.9 48 .18 638.6 0.09618 343.5 60 .52 917.2 0.1663 563.5 72 .55 1232 0.2385 775.4 84 .28 1588 0.3153 992.8 (HSOoF /nitroben zene solution) 0.3941 1205 ^0 .0000 IO.64 0 .06478 15.12 HP02F2(microburette) 0 .1920 25.68 uncorrected 0 .4311 33-44 0.000 9.274 0 .8207 31.99 3.847 24.92 1 .441 50.89 7.617 33.49 5 .550 114.7 13.31 41.57 16 .68 205.9 19.08 43.45 36 .95 480.7 26.77 40.34 81 .06 1504 34.85 36.50 42.16 34.73 insoluble TABLE 14 (cont'd)
H-SO, (solution) H?SO, (microburette)
2 2 10 m 107X 10 m 107X ohm- cm. ohm-lcmrl
0.000 • 8.326 0.000 3.302 0.05473 10.97 5.450 292.9 0.1375 16.92 8.022 359.8 . 0.1899 26.47 10.57 416.1 0.3293 39.66 15.49 504.7 0.5175 57.43 20.98 593.5 0.7729 81.84 29.26 698.9 1.410 105.0 39.39 800.8 2.347 152.0 51.49 832.3 5.361 231.5 65.84 -944.5 9.847 343.1 82.41 988.7 18.00 537.5 97.75 1062 34.02 1011 . Q 43.85 1376 EF. 0.000 5.964 0.000- 3.71 0.7123 40.72 27.5 914 1.893 77.07 46.2 1650 2.938 108.1 58.4 2950 4.676 155.4 7.499 238.3 11.60 369.4 18.28 655.9 28.28 812.2 46.72 1032 63.67 1210 87.57 1516 Addition of 0.0277 g-of water (0.0267 molal) causedX to approximately double Fig. 29 Specific Conductivities of Some Solutes in Nitrobenzene at 25(
1 . 025OCH I to B o
2000+
tetra n-butylammonium bromide
-4- 500
.10 2 20 30 uo 10 x molality, m be estimated on a purely electrostatic basis. The electrical
free energy for a pair of ions of charge +e and -e separated by
a distance r in a medium of dielectric constant €, is given by,
2 F - e /re o Assuming a value of 2A for r and values of € equal to 120 and
34-5 for €HSQ F and 6^ ^ ^ respectively, the free energy 3 6 5 2 difference,AF, for the two ions in the different media is 3400 cal./mole. On this basis the ratio of the equilibrium constants in the two media would be
-AF/RT "
K2
The dielectric constants used are the bulk dielectric constants for the liquids and are strictly an approximation for the values in the immediate vicinity of the ions. It seems likely that the bulk dielectric constant of HSO3F is enhanced greatly by the associated nature of the liquid and hence is probably closer to that of nitrobenzene in the vicinity of the ions. This would 2 mean that the value 3 x 10 calculated above is too large and thus it may be concluded that the large factor of 10? actually observed for the ratio of the equilibrium constants is only partially accounted for by the differences in the dielectric constants of the two rnedia>. More important is the fact that the
SO-^F ion is stabilized in HSO^F acid solutions through strong hydrogen bonding to the solvent molecules. Such stabilization is not possible in the non-protonic solvent, nitrobenzene.
B) Sulphuric acid solutions in nitrobenzene
The cryoscopic values (Table 13) show that less than 118
TABLE 15
Equivalent Conductivities
Tetra-n-butj'-laramonium H2S04 bromide 66 102m 102C IOTK. >A 102C ; TV 0.4170 0.500 50 1.0 (0.4 ) 0.005260 32.88 0.8340 1.00 70 0.70(0.2 ) 0.01284 32.44 1.668 2.00 90 0.45(0.2 ) 0.03268 31.67 4.170 5.00 180 O.38 (0.1 ) 30.39 8.340 10.0 315 0.32 (0.1 ) O.O8I84 28.24 16.68 20.0 540 0.26 (0.03) 0.2142 25.12 20.85 25.0 620 0.24 (0.03) 0.5515 , mo les/litre 30.0 720 0.23 (0.03) K=l62xl0~4 25.02 40.0 855 0.21 (0.03) Trimethylhydroxyammoniurn 33.36 50.0 41.70 965 0.19 (0.03) picrate 65 HSO^F 0.003178 17.19 0.007591 12.64 0. 4170 0.500 20 0.01802 8.960 0. 8340 1.00 30 0.20.405 (0.12 ) 0.03760 6.554 668 2.00 40 0.20 (0.1) 0.08016 4.688 41. 170 5.00 65 0.12 (0.04) 0.1750 3.292 8. 340 10.0 100 0.10 (0.04) 0.4329 2.17' 68 20.0 170 0.085(0.02) K=0.17xl0~4 moles/litr 2016. 85 25.0 210 0.081(0.01) 25. 02 30.0 275 0.080(0.007) Lithium picrate 66 33. 36 40.0 390 0.093'( 0.007) 41. 70 50.0 525 0.11 0.1091 0.1574 0.2671 O.I367 0.5754 0.1237 Tetra-n-butylammonium 0.7782 0.1132 bromide 0.9362 0.1080 2 1.074 0.0994 K-0.0006x10"*- moles/litre .10 C. vA 0.05969 30.64 0.1153 29.79 0.1993 28.27 0.2858 27.13 0.3779 26 0.4723 25.5.207 Fig.- 30 Equivalent Conductivities of Some Electrolytes in| Nitrobenzene at 25°
o
O H2S0^ • HSO^F < A lithium picrat .21 o •P O trimethylhy- •H droxy-ammonium > picrate •H -P O 0
O O
C CO
10 C: one mole of particles are produced per mole of sulphuric acid molecules over the concentration range studied. The H^SO^ curve in Fig. 28 has a slope of 4-66 which is considerably less than that expected for a non-electrolyte and it would appear that some dimerization of the H^SO^ molecules occurs. An examination of the chemical and physical properties of the acids, H^SO^ and HSO^F, indicates that sulphuric acid is much more extensively hydrogen bonded (i.e. associated) than fluorosulphuric acid and it is thus not surprising that this tendency towards association is exhibited 'to a larger degree by HgSO^ in nitrobenzene. As was found in the case of the fluorosulphuric acid solutions the reproducibility of the conductivity values was poor (Table 11+). It is immediately apparent that, compared to the fairly strong electrolyte tetrabutylammonium bromide, J^SO^ like HSO^F appears to be a very weak electrolyte. The interpretation of the con• ductivity data for sulphuric acid is complicated by the fact that the acid itself undergoes rather extensive and complicated self dissociation reactions and it is difficult to predict what effect the nitrobenzene will have on these. Nevertheless, it is clear from the estimated equivalent conductivity values for the acid (Table 15) that the degree of dissociation into ions is not much greater than that observed for the fluorosulphuric acid solutions and it is reasonable to conclude that no significant dissociation according to
C6H5N02-HA ^ C6H5N02H+ + A" ^.?>) takes place. Ignoring the small degree of dissociation into ions the cryoscopic V values may be interpreted in terms of the reactio (H2S04)2 ^ 2H2S04 with an equilibrium constant of approximately 1.5 x 10"^ mole kg.
It is evident that in solutions of nitrobenzene the acid strength of sulphuric acid relative to the base nitrobenzene has been reduced not only to the extent that there is no significant protonation of nitrobenzene but also (as is shown by the observa• tion that approximately half of the sulphuric acid molecules are present as dimers) to the extent that the nitrobenzene is in com• petition with other molecules of H2SO4 for forming hydrogen- bonded adducts. An earlier report that sulphuric acid is com• pletely dimerized in nitrobenzene solutions was not confirmed in this work.
C) Difluorophosphoric acid solutions
In the case of HP0pF2, the acid and nitrobenzene were found to be immiscible and only at the lowest concentrations did dissolution appear to be complete (at a concentration of less -
than 0.1 molal), whereas H2S0^ and HSO^F dissolve readily. The electrical conductivities observed for the HPOgFg/nitrobenzene solutions were small and quickly levelled off at a value of -7 1-1 43 x 10 ' ohms "cm. . The data for the conductivity studies for
HF/nitrobenzene solutions obtained by other workers cite also given in Table 14 and it is interesting to note that at low concentra• tions the specific conductivity is of the same order as that for the inorganic acids studied in this work. 5.4. Conclusion
The acid strengths of HSO3F and H2S0^ defined in terms of their ability to protonate bases depends to a large degree on the nature of the solvent. In bulk, where the acids themselves provide the protonic medium, or in solution in other protonic solvents these acids are indeed strong; however, their strengths may be considerably reduced in media where there is no- stabili• zation by hydrogen bonding to solvent molecules of the anion pro• duced by the protonation reaction. The importance of anion stabilization by hydrogen bonding to solvent molecules in deter• mining the apparent strengths of acids was recently mentioned
71 elsewhere.'
Although HSO3F and HgSO^ do not exhibit acid behaviour in dilute solutions in nitrobenzene in the sense of protonating nitrobenzene molecules the fact that the acids are soluble in• dicates rather weak, acid behaviour of these solutes towards nitrobenzene. In order for these acids to be soluble their ex• tensive hydrogen bonded structures must be destroyed and this, presumably, is due to interaction between the acid molecules and the nitrobenzene molecules in the form of hydrogen bonding between the two (a type of weak acid-base reaction). The fact that HPO2F2 is essentially insoluble in nitrobenzene suggests that HPO2F2 is not able to form strong hydrogen bonds with nitrobenzene molecules and thus indicates that even in non-protonic media difluorophos•
phoric acid is a weaker acid than either HoS0, or HS0oF. 123
CHAPTER VI
Summary and Suggestions for Further Work
6.1 Summary
From the investigation of the electrical conductivities of solutions of the alkali and alkaline earth metal difluoro- phosphates in HPO2F2 ^w0 models were proposed to explain the results: (i) complete, dissociation with weak solute-solvent interaction and (ii) incomplete dissociation with the lithium salt exhibiting the greatest degree of dissociation of the alkali metal difluorophosphates. Of the two models, the latter is pre• ferred. Nuclear magnetic resonance studies of alkali metal di- fluorophosphate solutions tend to substantiate this preference, especially the 1°F chemical shifts which are small and to low-field. It is probable that if difluorophosphates are fully dissociated then the presence of a large concentration of POgFg" ions would result in chemical shifts to high-field due to in• creased shielding of the fluorine nuclei.
Various bases were also studied conductimetrically in
HPO2F2. Organic amines gave conductivities of the same order as
that exhibited by CsP02F2 solutions, thus suggesting that complete dissociation of the amines does not occur. Several of the organic compounds investigated proved to be virtually insoluble. In• organic bases were also examined; the alkali and alkaline earth metal chlorides gave conductivities similar to those given by the corresponding difluorophosphates thus suggesting that the chlorides undergo complete solvolysis and the HC1 produced is a non-electrolyte in HPO2F2. Of the chlorides studied tetraphenylarsonium chloride gave the highest conductivity which is indicative of large cations exhibiting the least' ion-pairing and greatest conductivity in a medium of low dielectric constant.
Solutions of SbF^, HSO^F and HgSO^ were investigated conductimetrically in HPO2F2 and it was concluded that these solutes are acids and can be titrated with base. In each case the salt produced by the titration was insoluble and was isolated.
The insolubility of these salts e.g. KSbF^P02F2, KSO^F and KHSO^ lends further support to the proposal that HPO2F2 has a low di• electric constant. Nuclear magnetic resonance studies of the
SbF^/HP02F2 solutions were also examined and it was found that
the expected species, HSbFr)P02F2 could be identified in the various spectra; other antimony pentafluoride—difluorophosphate com• plexes are proposed to explain the remaining peaks in the spectra.
An attempt was.made to identify further the species SbF^POgFg"
by decomposing KSbF^P02F2, however, of the expected products,
PO2F and KSbF^, only the hexafluoroantimonate was found.
Solutions of HgSO^, HP02F2 and HSO^F were also studied in nitrobenzene to investigate further the factors affecting relative acid strengths of protonic acids. Electrical conduc• tivity studies showed that HgSO^ and HSO^F were soluble and were
virtually non-electrolytes whereas HP02F2 was only slightly soluble. Cryoscopic investigations of the HgSO^ and HSO^F solutions in nitrobenzene indicate that HSO^F exists essentially
as a monomer whereas H2S0, exists in a polymeric form. 6.2 Suggestions for further work
As has been mentioned previously, a knowledge of the di• electric constant of HPC^Fg would be of great value in the inter- 72 pretation of the conductivity data. Gillespie et al. have re• ported on the problems of measuring the dielectric constant of conducting and viscous liquids e.g. sulphuric, nitric and chloro- sulphuric acids. The "force" method of determining the static dielectric constant was used and unfortunately apparatus of this type is not available in these laboratories. It also will be necessary to undertake the' measurement of transport numbers to determine whether abnormal conductivity occurs in the solvent and to what extent ions other than the autoprotolysis ions carry the 55 current. However, if the method used by Thompson ' is employed, the analytical problems of measuring small changes in concentra• tion in the anode and cathode compartments of the transport number cell will have to be solved.
Cryoscopy has proved very useful in the investigation of
^2^4 and HSO-jF solutions when used in conjunction with conduc• tivity. Information regarding degrees of dissociation of various electrolytes in these media can be obtained. The cryoscopic constant for HPOgFg is not known and would have to be determined, otherwise only relative extents of dissociation could be found but this would still be useful in the interpretation of the con• ductivity results of alkali metal difluorophosphates.
is a • It has been confirmed that HP02F2 weaker acid than both H0SO1 and HSOoF and-it would be interesting to attempt a 126
determination of the Hammet acidity function Ho for HPOgFg. It
is likely that AsF^ and BF^ will also prove to be acids in HPO2F2
and titration of these acids with KPO2F2 may yield insoluble salts
of composition KBF3P02F2 and KAsFr_P02F2. As alkaline earth and
alkali metal difluorophosphates have now been prepared it would
be of interest to attempt the preparation of transition metal
difluorophosphates by reaction of the transition metal chlorides with difluorophosphoric acid. 127
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