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Thermodynamics of formation, hydrolysis and ionization of monofluorophosphorous acid
Po/y/,edron Vol. 10, No. 14, pp. 169W698, Printed in Great Britain
1991 0
0277%5387/91 $3.00+.00 I991 Pergamon Press plc
THERMODYNAMICS OF FORMATION, HYDROLYSIS AND IONIZATION OF MONOFLUOROPHOSPHOROUS ACID JOHN
W. LARSON
Department of Chemistry, Marshall University, Huntington, WV 25755-2520, U.S.A. (Received 22 October 1990 ; accepted 19 April 1991)
Abstract-A NMR study of the equilibrium reaction of HF with aqueous phosphorous acid between 8 and 65°C and various hydrogen ions concentrations gave values of AG% = 2.2 kJ, AHe = -26 kJ and AS0 = - 94 J K- I for the acid ionization of monofluorophosphorous acid and AGe = - 16.2 kJ, AHe = - 16.9 kJ and ASe = -2.2 J K- ’ for the hydrolysis of monofluorophosphorous acid to phosphorous acid and HF.
Monofluorophosphorous acid, H[FHPO,], was first synthesized in 1968 by Centofanti and Parry’ from the disproportionation of F*HPO. Several additional preparative routes have subsequently been discovered.*-’ Other names used for this acid are phosphinic fluoride, phosphonofluoridous acid and phosphonofluoridic acid. In this work, aqueous FHPOzH is shown to be formed by a fast equilibrium reaction between aqueous HF and HP03H2, e.g. (1) :
H20( 1) + FHPO,H(aq).
(1)
Thermodynamic properties of monofluorophosphorous acid were obtained from the study of the temperature and hydrogen ion concentration dependence of the equilibrium quotient of this reaction. The thermochemical properties of FHP02H are of interest because : (i) FHP02H contains only four heavy atoms and therefore this work may supply useful reference data in relation to theoretical calculations on important phosphorus compounds and reactions, particularly, hydrolysis reactions. (ii) This data may further our understanding of the effects of substituents that are bonded directly to second row atoms. (iii) The thermodynamics of hydrolysis of FHPO*H [reverse reaction (l)] may be related to and could be helpful in the understanding of biologically important phosphate hydrolysis reactions. No thermochemical properties of this compound are currently available.
EXPERIMENTAL A concentration solution of phosphorous acid (5.28 mmol of H3P03 per gram of solution) was made from the pure acid obtained from Aldrich. The concentration of this solution and concentrated solutions of hydrofluoric acid (19.85 mmol g- ‘) and concentrated hydrochloric acid (10.22 mmol g- I), both obtained from Fischer Scientific, were analysed by standard acidimetric titration methods. The HF was stored, dispensed and analysed using plastic labware. The sodium phosphite pentahydrate (BDH Chemicals) was analysed for water by heating in vacua to 200°C. The sample solutions used in this work were prepared by weight with concentrations expressed in terms of molalities and the standard state for the solute taken as the hypothetical one molal solution. Phosphorus magnetic resonance spectra were obtained using a Varian XL-200 spectrometer. All spectra involving HF solutions were recorded in 10 mm NMR tubes with Teflon sleeves and caps. Temperatures were controlled to within f0.3”C and calibrated with a mercury thermometer. The sensitivity of the NMR to low concentration phosphorus species, under the conditions of the equilibrium measurements, was investigated by making up standard solutions containing 0.5 m Na2HP03 and small amounts of NaH2P02. The detection limit of phosphorus as H2P02- was found to be about 0.0003 m or 0.0006 mole fraction. The equilibrium fractions of phosphorus present as FHPO*H in the present work covered the range
1695
1696
J. W. LARSON
from 0.013 to 0.007. The intrinsic error in equilibrium quotients is therefore estimated to be between 5 and 10%. This is comparable with the actual scatter in the experimental data. CALCULATIONS
AND RESULTS
When hydrofluoric acid is added to a solution of phosphorous acid, small peaks emerge immediately corresponding to the doublet of doublets of FHPO*H formed by reaction (1). For the whole temperature range between 5 and 85°C reaction (1) appears to be fast on the time scale it takes to make the solutions and record the spectrum (20 min) and slow on the NMR time scale that would lead to line broadening (0.01 s). At high HF and H+ concentrations and at high temperatures, the FHPOIH peaks broaden appreciably and eventually the PF splitting disappears. This is apparently due to FHPO,H exchanging F with the HF in solution. At the concentrations of the equilibrium measurements reported here, the broadening is slight and has no effect on the measurements. Hydrogen decoupling collapses the FHPO*H spectrum to a doublet. The coupling constants and chemical shifts of the aqueous acid and its conjugate base are reported in Table 1 together with the corresponding value previously determined for liquid monofluorophosphorous acid and aqueous phosphorous acid. The equilibrium expression for reaction (1) may be written as eq. (2) : K, =
mFHm,H
(2)
mHFmHPO,H,
In this equation and all subsequent equations, the activity coefficients of the neutral species are estimated to be unity and are left out of the equilibrium expressions. The fraction of the phosphorous that is fluorinated as either FHPO,H or FHP02-, XPF, is determined directly from integration of the NMR
Table
peaks. Equation (2) may therefore be re-written as eq. (3) K, =
m-d
1- &)
where Q is the equilibrium quotient for the reaction and XI and X, are the respective fractions of phosphite and fluorophosphite in the acid form. Both of these fractions may be obtained via the measured P-H coupling constant of phosphite. Xi is calculated directly from the measured coupling constant by the method given in ref. 6, eq. (4) : X, = (J-629.5)/56.5.
JPHU-W JPF (Hz) 6, (ppm)
712 987 4.0
L?Reference 1. bReference 6. ‘Values are apparently only f 5 Hz.
FHPO,-(aq) 698’ 973’ 4.0
(4)
The change in the coupling constants of FHP02 when it undergoes protonation is too small and the measured values are too uncertain to apply the same procedure in calculating X,. This fraction is instead obtained by considering the acid equilibrium reaction, eq. (5), and the equilibrium expression, eq. (6), where y is the mean activity coefficient of H+ and FHPO*-: FHPO,H(aq)
e
H+(aq) + FHPO,-(aq)
(5)
KS = mhmFHP0,y2/mFHP0,n= H(l --XZ)/~Z. (6) The PH coupling constant of phosphorous acid is used a second time in an empirical method of obtaining the value of mny’, defined now as H, from the known acid ionization constant for phosphorous acid and the value of 11, eq. (7) : H = mHy2 =
{Ka,~3~,)~~/(l--XI).
(7)
Application of eq. (7) to find the value of mHy2 appropriate for eq. (6) assumes that the activity coefficient of FHP02- is approximately equal to HP03H-. Considering the similar structures of these species, this assumption is likely to be quite accurate. The fraction of FHP02H in the acid form is then given by eq. (8) : X2 = l/(1 + KS/H).
1. Coupling constants and chemical shifts (relative to 85% phosphoric monofluorophosphorous and phosphorous acid and their anions FHP02H(aq)
(3)
x1
acid) for
FHPO,H(l)”
HPO&I,(aq)*
783 1028 -2.74
686
629.5
5.88
3.22
effected by hydrogen-bonding
HPO,H-(aq)b
to HF and are therefore known to
(8)
Monofluorophosphorous
1697
acid
F- with HF,7 and function H has been defined previously. The uncertainties given in Table 2 are the total uncertainties including the uncertainty involved in the estimates of activity coefficients, the uncertainty in the reference values of K, of phosphorous acid, and the uncertainty involved in the measurement of X,,.
O/Xl
i/H
Fig. 1. Determination of K, and K, by graphing Q/X, vs l/H, eq. (9) at 65°C.
Substituting eq. (8) for X2 in eq. (3) and rearranging results in eq. (9) :
Q/xl = KI +K,fMlIH).
(9)
Use of this equation to obtain K, and KS is demonstrated in Fig. 1. Measurements were made on five different solutions with H ranging between 0.05 and 1.5 m at temperatures of 8, 25,45 and 65°C to obtain the values given in Table 2. The concentration ranges in these measurements were: 1.754.2 m for HF; 0.15-1.3 m for HP03H2; O-l.3 m for HCl ; and c2.4 for Na2HP03. None of these concentrations enter directly into the calculations of the equilibrium constants except for the concentration of HF. In the calculation of Q, the initial concentration of HF must be corrected for the amount that reacts to form F-, FHPO*H, FHPO*and FHF-. Under the conditions of our measurements, calculations indicate less than 2% of the HF is present in the first three forms. At the highest pH, the concentration of FHF- becomes appreciable, about 20%. The amount of FHF- present was easily calculated, eq. (10) : mFHFm
=
m%G&IH~
(10)
where K, is the acid dissociation constant of HF,’ Kd is the equilibrium constant of the reaction of
Thermodynamics of ionization of monojluorophosphorous acid
The van? Hoff isochore graph of the equilibrium constants is shown in Fig. 2. Standard thermodynamic analysis leads to values of AGF = +2.2fl.l kJ, AH? = -26.0f4.0 kJ and ASP = - 94 f 13 J K- ’ for the acid ionization process. FHP02H is a stronger acid than H2P02H by 0.87 pK, units and is the strongest known substituted phosphinic acid* of formula XHP02H. This fluorosubstituent effect is somewhat larger than the 0.6 units difference between the pKal values of FP03H2 and HP03H2 and smaller than the 1.6 units found between their pK, values. The acidity difference appears to be due entirely to the enthalpy change with the fluoro-substituted acid having a 10 kJ more exothermic heat of ionization. Thermodynamics of hydrolysis of monojkorophosphorous acid
The value of AH? = 16.9 f 1.3 kJ obtained from the slope of In K, against l/T is combined with the AG? = 16.2kO.6 kJ at 25°C to obtain AS? = 2.224.3 J K- ‘. The small entropy change is reasonable considering the symmetric nature of reaction (1) and the fact that all of the reactants and products are neutral. Literature values’,” for the thermodynamics of HF(aq), H,PO,(aq) and H,O(l) have been used together with the thermodynamics of reaction (1)
0
______~
Table 2. Equilibrium constants for reactions (1) and (4) and the pK, of monofluorophosphorous acid ,“(K)++y
Temperature (“C)
103K,”
KS
pKab
8 25 45 65
0.95 1.4 2.0 3.4
0.71 0.43 0.27 0.10
0.15 0.37 0.56 1.0
“Uncertainty ‘Uncertainty
= f 30%. = +0.05+ 10%.
In(K1) -,I / -8 L__ 2.8
ic__b I _e 2.9
3
3.1
3.2 1000/T
3.3
3.4
3.5
1 3.8
Fig. 2. The van? Hoff isochore graphs for reactions (1) and (5).
1698
J. W. LARSON
Table 3. Thermodynamics of formation’ of monofluorophosphorous, monofluorophosphite, phosphorous and phosphite at 25°C AGrf, (kJ) HPO,H,(aq) HPO,H-(aq) FHPO,H(aq) FHPO,-(aq)
AZZre (kJ)
- 836.6 - 829.4 -880.1 - 877.9
-954.4 (- 969.4)+ -971.0 -997.8
S* (149)” 75 170 76
a Reference values in parentheses. bReference 9. ‘Based on ?&,, from ref. 9 and Snlpo,-SH~po~ = 8.8 from ref. 10.
and the thermodynamics of ionization of phosphorous and monofluorophosphorous acid to calculate the thermodynamics of formation reported in Table 3. Phosphorus appears to be at the borderline between the more electropositive elements which form fluorides that are stable in water and the more electronegative elements which form fluorides that hydrolyse completely in water. Thus SiF,‘- is thermodynamically stable in aqueous solutions, while S03F- is readily hydrolysed. In strong acid solutions, the enthalpy of hydrolysis of these species is calculated to be + 24 and - 50.1 kJ per mol of HF,
respectively. The enthalpy of hydrolysis of FHP02 H(aq), - 16.9 kJ, is intermediate between these values. Related values calculated from literature data’ are : PFs-(aq)+4H20(1)+ H2P04-(aq)+6HF(aq)
AH = - 19.6 kJ
HPOK(aq)+H,O(l)e H,PO,-(aq)
+ HF(aq)
AG =
+ 8.2 kJ.
The calculated value of AG for this second reaction seems peculiar in light of the results of this work on FHP02H and the tendency of fluorophosphate to undergo hydrolysis. The thermochemistry of fluorophosphoric acid is therefore currently being re-investigated along with additional substituted phosphonic and phosphinic acids.
REFERENCES 1. L. F. Centofanti and R. W. Parry, Znorg. Chem. 1968, 7, 1005. 2. G. A. Olah and C. W. McFarland, Znorg. Chem. 1972, 11, 845.
3. R. Bender, C. Demay, J. C. Elkaim and J. G. Reiss, Phosphorus 1974,4, 183.
4. H. Falius, Angew. Chem., Znt. Edn Engl. 1970, 9, 733.
5. B. Blazer and K. H. Worms, Anorg. Allg. Chem. 1968,370,
117.
6. J. W. Larson and M. Pippin, Polyhedron 1989, 8, 527.
7. L. G. Sillen and A. E. Martell, Stability Constants of Metal Zon Complexes. The Chemical Society, London (1964). 8. W. P. Jencks and J. Regenstein, Handbook of Biochemistry and Molecular Biology (Edited by G. Fasman), Vol. 1, 3rd edn. CRC Press, Cleveland (1976). 9. D. D. Wagman, W. H. Evans, V. B. Parker, R. H. Schumm, I. Halow, S. M. Bailey, K. L. Chumey and R. L. Nuttall, J. Phys. Chem. Ref: Data 1982, 11, 1. 10. W. M. Latimer, Oxidation Potentials. Prentice-Hall, New York (1952).
1991 0
0277%5387/91 $3.00+.00 I991 Pergamon Press plc
THERMODYNAMICS OF FORMATION, HYDROLYSIS AND IONIZATION OF MONOFLUOROPHOSPHOROUS ACID JOHN
W. LARSON
Department of Chemistry, Marshall University, Huntington, WV 25755-2520, U.S.A. (Received 22 October 1990 ; accepted 19 April 1991)
Abstract-A NMR study of the equilibrium reaction of HF with aqueous phosphorous acid between 8 and 65°C and various hydrogen ions concentrations gave values of AG% = 2.2 kJ, AHe = -26 kJ and AS0 = - 94 J K- I for the acid ionization of monofluorophosphorous acid and AGe = - 16.2 kJ, AHe = - 16.9 kJ and ASe = -2.2 J K- ’ for the hydrolysis of monofluorophosphorous acid to phosphorous acid and HF.
Monofluorophosphorous acid, H[FHPO,], was first synthesized in 1968 by Centofanti and Parry’ from the disproportionation of F*HPO. Several additional preparative routes have subsequently been discovered.*-’ Other names used for this acid are phosphinic fluoride, phosphonofluoridous acid and phosphonofluoridic acid. In this work, aqueous FHPOzH is shown to be formed by a fast equilibrium reaction between aqueous HF and HP03H2, e.g. (1) :
H20( 1) + FHPO,H(aq).
(1)
Thermodynamic properties of monofluorophosphorous acid were obtained from the study of the temperature and hydrogen ion concentration dependence of the equilibrium quotient of this reaction. The thermochemical properties of FHP02H are of interest because : (i) FHP02H contains only four heavy atoms and therefore this work may supply useful reference data in relation to theoretical calculations on important phosphorus compounds and reactions, particularly, hydrolysis reactions. (ii) This data may further our understanding of the effects of substituents that are bonded directly to second row atoms. (iii) The thermodynamics of hydrolysis of FHPO*H [reverse reaction (l)] may be related to and could be helpful in the understanding of biologically important phosphate hydrolysis reactions. No thermochemical properties of this compound are currently available.
EXPERIMENTAL A concentration solution of phosphorous acid (5.28 mmol of H3P03 per gram of solution) was made from the pure acid obtained from Aldrich. The concentration of this solution and concentrated solutions of hydrofluoric acid (19.85 mmol g- ‘) and concentrated hydrochloric acid (10.22 mmol g- I), both obtained from Fischer Scientific, were analysed by standard acidimetric titration methods. The HF was stored, dispensed and analysed using plastic labware. The sodium phosphite pentahydrate (BDH Chemicals) was analysed for water by heating in vacua to 200°C. The sample solutions used in this work were prepared by weight with concentrations expressed in terms of molalities and the standard state for the solute taken as the hypothetical one molal solution. Phosphorus magnetic resonance spectra were obtained using a Varian XL-200 spectrometer. All spectra involving HF solutions were recorded in 10 mm NMR tubes with Teflon sleeves and caps. Temperatures were controlled to within f0.3”C and calibrated with a mercury thermometer. The sensitivity of the NMR to low concentration phosphorus species, under the conditions of the equilibrium measurements, was investigated by making up standard solutions containing 0.5 m Na2HP03 and small amounts of NaH2P02. The detection limit of phosphorus as H2P02- was found to be about 0.0003 m or 0.0006 mole fraction. The equilibrium fractions of phosphorus present as FHPO*H in the present work covered the range
1695
1696
J. W. LARSON
from 0.013 to 0.007. The intrinsic error in equilibrium quotients is therefore estimated to be between 5 and 10%. This is comparable with the actual scatter in the experimental data. CALCULATIONS
AND RESULTS
When hydrofluoric acid is added to a solution of phosphorous acid, small peaks emerge immediately corresponding to the doublet of doublets of FHPO*H formed by reaction (1). For the whole temperature range between 5 and 85°C reaction (1) appears to be fast on the time scale it takes to make the solutions and record the spectrum (20 min) and slow on the NMR time scale that would lead to line broadening (0.01 s). At high HF and H+ concentrations and at high temperatures, the FHPOIH peaks broaden appreciably and eventually the PF splitting disappears. This is apparently due to FHPO,H exchanging F with the HF in solution. At the concentrations of the equilibrium measurements reported here, the broadening is slight and has no effect on the measurements. Hydrogen decoupling collapses the FHPO*H spectrum to a doublet. The coupling constants and chemical shifts of the aqueous acid and its conjugate base are reported in Table 1 together with the corresponding value previously determined for liquid monofluorophosphorous acid and aqueous phosphorous acid. The equilibrium expression for reaction (1) may be written as eq. (2) : K, =
mFHm,H
(2)
mHFmHPO,H,
In this equation and all subsequent equations, the activity coefficients of the neutral species are estimated to be unity and are left out of the equilibrium expressions. The fraction of the phosphorous that is fluorinated as either FHPO,H or FHP02-, XPF, is determined directly from integration of the NMR
Table
peaks. Equation (2) may therefore be re-written as eq. (3) K, =
m-d
1- &)
where Q is the equilibrium quotient for the reaction and XI and X, are the respective fractions of phosphite and fluorophosphite in the acid form. Both of these fractions may be obtained via the measured P-H coupling constant of phosphite. Xi is calculated directly from the measured coupling constant by the method given in ref. 6, eq. (4) : X, = (J-629.5)/56.5.
JPHU-W JPF (Hz) 6, (ppm)
712 987 4.0
L?Reference 1. bReference 6. ‘Values are apparently only f 5 Hz.
FHPO,-(aq) 698’ 973’ 4.0
(4)
The change in the coupling constants of FHP02 when it undergoes protonation is too small and the measured values are too uncertain to apply the same procedure in calculating X,. This fraction is instead obtained by considering the acid equilibrium reaction, eq. (5), and the equilibrium expression, eq. (6), where y is the mean activity coefficient of H+ and FHPO*-: FHPO,H(aq)
e
H+(aq) + FHPO,-(aq)
(5)
KS = mhmFHP0,y2/mFHP0,n= H(l --XZ)/~Z. (6) The PH coupling constant of phosphorous acid is used a second time in an empirical method of obtaining the value of mny’, defined now as H, from the known acid ionization constant for phosphorous acid and the value of 11, eq. (7) : H = mHy2 =
{Ka,~3~,)~~/(l--XI).
(7)
Application of eq. (7) to find the value of mHy2 appropriate for eq. (6) assumes that the activity coefficient of FHP02- is approximately equal to HP03H-. Considering the similar structures of these species, this assumption is likely to be quite accurate. The fraction of FHP02H in the acid form is then given by eq. (8) : X2 = l/(1 + KS/H).
1. Coupling constants and chemical shifts (relative to 85% phosphoric monofluorophosphorous and phosphorous acid and their anions FHP02H(aq)
(3)
x1
acid) for
FHPO,H(l)”
HPO&I,(aq)*
783 1028 -2.74
686
629.5
5.88
3.22
effected by hydrogen-bonding
HPO,H-(aq)b
to HF and are therefore known to
(8)
Monofluorophosphorous
1697
acid
F- with HF,7 and function H has been defined previously. The uncertainties given in Table 2 are the total uncertainties including the uncertainty involved in the estimates of activity coefficients, the uncertainty in the reference values of K, of phosphorous acid, and the uncertainty involved in the measurement of X,,.
O/Xl
i/H
Fig. 1. Determination of K, and K, by graphing Q/X, vs l/H, eq. (9) at 65°C.
Substituting eq. (8) for X2 in eq. (3) and rearranging results in eq. (9) :
Q/xl = KI +K,fMlIH).
(9)
Use of this equation to obtain K, and KS is demonstrated in Fig. 1. Measurements were made on five different solutions with H ranging between 0.05 and 1.5 m at temperatures of 8, 25,45 and 65°C to obtain the values given in Table 2. The concentration ranges in these measurements were: 1.754.2 m for HF; 0.15-1.3 m for HP03H2; O-l.3 m for HCl ; and c2.4 for Na2HP03. None of these concentrations enter directly into the calculations of the equilibrium constants except for the concentration of HF. In the calculation of Q, the initial concentration of HF must be corrected for the amount that reacts to form F-, FHPO*H, FHPO*and FHF-. Under the conditions of our measurements, calculations indicate less than 2% of the HF is present in the first three forms. At the highest pH, the concentration of FHF- becomes appreciable, about 20%. The amount of FHF- present was easily calculated, eq. (10) : mFHFm
=
m%G&IH~
(10)
where K, is the acid dissociation constant of HF,’ Kd is the equilibrium constant of the reaction of
Thermodynamics of ionization of monojluorophosphorous acid
The van? Hoff isochore graph of the equilibrium constants is shown in Fig. 2. Standard thermodynamic analysis leads to values of AGF = +2.2fl.l kJ, AH? = -26.0f4.0 kJ and ASP = - 94 f 13 J K- ’ for the acid ionization process. FHP02H is a stronger acid than H2P02H by 0.87 pK, units and is the strongest known substituted phosphinic acid* of formula XHP02H. This fluorosubstituent effect is somewhat larger than the 0.6 units difference between the pKal values of FP03H2 and HP03H2 and smaller than the 1.6 units found between their pK, values. The acidity difference appears to be due entirely to the enthalpy change with the fluoro-substituted acid having a 10 kJ more exothermic heat of ionization. Thermodynamics of hydrolysis of monojkorophosphorous acid
The value of AH? = 16.9 f 1.3 kJ obtained from the slope of In K, against l/T is combined with the AG? = 16.2kO.6 kJ at 25°C to obtain AS? = 2.224.3 J K- ‘. The small entropy change is reasonable considering the symmetric nature of reaction (1) and the fact that all of the reactants and products are neutral. Literature values’,” for the thermodynamics of HF(aq), H,PO,(aq) and H,O(l) have been used together with the thermodynamics of reaction (1)
0
______~
Table 2. Equilibrium constants for reactions (1) and (4) and the pK, of monofluorophosphorous acid ,“(K)++y
Temperature (“C)
103K,”
KS
pKab
8 25 45 65
0.95 1.4 2.0 3.4
0.71 0.43 0.27 0.10
0.15 0.37 0.56 1.0
“Uncertainty ‘Uncertainty
= f 30%. = +0.05+ 10%.
In(K1) -,I / -8 L__ 2.8
ic__b I _e 2.9
3
3.1
3.2 1000/T
3.3
3.4
3.5
1 3.8
Fig. 2. The van? Hoff isochore graphs for reactions (1) and (5).
1698
J. W. LARSON
Table 3. Thermodynamics of formation’ of monofluorophosphorous, monofluorophosphite, phosphorous and phosphite at 25°C AGrf, (kJ) HPO,H,(aq) HPO,H-(aq) FHPO,H(aq) FHPO,-(aq)
AZZre (kJ)
- 836.6 - 829.4 -880.1 - 877.9
-954.4 (- 969.4)+ -971.0 -997.8
S* (149)” 75 170 76
a Reference values in parentheses. bReference 9. ‘Based on ?&,, from ref. 9 and Snlpo,-SH~po~ = 8.8 from ref. 10.
and the thermodynamics of ionization of phosphorous and monofluorophosphorous acid to calculate the thermodynamics of formation reported in Table 3. Phosphorus appears to be at the borderline between the more electropositive elements which form fluorides that are stable in water and the more electronegative elements which form fluorides that hydrolyse completely in water. Thus SiF,‘- is thermodynamically stable in aqueous solutions, while S03F- is readily hydrolysed. In strong acid solutions, the enthalpy of hydrolysis of these species is calculated to be + 24 and - 50.1 kJ per mol of HF,
respectively. The enthalpy of hydrolysis of FHP02 H(aq), - 16.9 kJ, is intermediate between these values. Related values calculated from literature data’ are : PFs-(aq)+4H20(1)+ H2P04-(aq)+6HF(aq)
AH = - 19.6 kJ
HPOK(aq)+H,O(l)e H,PO,-(aq)
+ HF(aq)
AG =
+ 8.2 kJ.
The calculated value of AG for this second reaction seems peculiar in light of the results of this work on FHP02H and the tendency of fluorophosphate to undergo hydrolysis. The thermochemistry of fluorophosphoric acid is therefore currently being re-investigated along with additional substituted phosphonic and phosphinic acids.
REFERENCES 1. L. F. Centofanti and R. W. Parry, Znorg. Chem. 1968, 7, 1005. 2. G. A. Olah and C. W. McFarland, Znorg. Chem. 1972, 11, 845.
3. R. Bender, C. Demay, J. C. Elkaim and J. G. Reiss, Phosphorus 1974,4, 183.
4. H. Falius, Angew. Chem., Znt. Edn Engl. 1970, 9, 733.
5. B. Blazer and K. H. Worms, Anorg. Allg. Chem. 1968,370,
117.
6. J. W. Larson and M. Pippin, Polyhedron 1989, 8, 527.
7. L. G. Sillen and A. E. Martell, Stability Constants of Metal Zon Complexes. The Chemical Society, London (1964). 8. W. P. Jencks and J. Regenstein, Handbook of Biochemistry and Molecular Biology (Edited by G. Fasman), Vol. 1, 3rd edn. CRC Press, Cleveland (1976). 9. D. D. Wagman, W. H. Evans, V. B. Parker, R. H. Schumm, I. Halow, S. M. Bailey, K. L. Chumey and R. L. Nuttall, J. Phys. Chem. Ref: Data 1982, 11, 1. 10. W. M. Latimer, Oxidation Potentials. Prentice-Hall, New York (1952).