J. Phys. Chem. 1996, 100, 2443-2446
2443
Free Energies of Hydration in the Gas Phase of Some Phosphate Singly and Doubly Charged Anions: (HO)2PO2- (Orthophosphate), (HO)O2POPO2(OH)2- (Diphosphate), Ribose 5-Phosphate, Adenosine 5′-Phosphate, and Adenosine 5′-Diphosphate Arthur T. Blades, Yeunghaw Ho, and Paul Kebarle* Department of Chemistry, UniVersity of Alberta, Edmonton, AB, Canada T6G 2G2 ReceiVed: July 12, 1995; In Final Form: NoVember 1, 1995X
Singly and doubly charged anions, AZ-, of phosphates, sugar phosphates, and adenosine phosphates are obtained in the gas phase by electrospray. The ions are introduced into a reaction chamber where sequential hydration equilibria (n - 1,n) AZ-(H2O)n-1 + H2O ) AZ-(H2O)n occur. The equilibrium constants Kn-1,n are determined with a mass spectrometer. Free energies ∆G°n-1,n at room temperature were determined for (HO)2PO2-, orthophosphate, (HO)O2POPO2(OH)22-, diphosphate, ribose 5-phosphate singly charged, adenosine 5′phosphate, and singly charged and doubly charged adenosine 5′-diphosphate. The magnitudes of the hydration energies provide information on the stability of the ions. Unusually low hydration exoergicities indicate the presence of intramolecular hydrogen bonding. Evidence is provided for the presence of intramolecular hydrogen bonding in ribose phosphate, adenosine monophosphate, and adenosine diphosphate.
Introduction Determinations of gas-phase equilibria involving singly charged ions and solvent molecules such as H2O (see eq 1) or
A-(H2O)n-1 + H2O ) A-(H2O)n
(n - 1, n)
(1)
other ligands, were initiated some 30 years ago.1,2 The sequential bond free energies, ∆G°n-1,n and enthalpies, ∆H°n-1,n, resulting from such equilibria measurements have provided a wealth of data2,3 on ion-solvent and ion-ligand interactions for positive M+ and negative A- ions. Many ions of great interest in condensed-phase chemistry and biochemistry could not be produced in the gas phase by the conventional methods.1-3 For example, negative ions A- are obtained by deprotonation of AH in the gas phase but this method requires that the AH compounds be sufficiently volatile. Electrospray mass spectrometry (ESMS) developed by John Fenn and co-workers4 is a method with which electrolyte ions present in solution can be transferred to the gas phase.5 With ES it is possible to obtain gas phase singly or multiply charged ion-ligand complexes,6 multiply protonated peptides and proteins,7 doubly charged anions8 such as SO42- and multiply deprotonated nucleic acids.9 Recently, we described10,11 an ion source reaction chamber with which ion-molecule equilibria involving ES produced ions can be determined. These first experiments10,11 involved positive ions. An investigation involving some 50 singly and doubly charged anions deriving from oxoacids of C, N, S, P, Cl, Br, and I was completed more recently.12 The present work focuses on the hydration of the orthophosphate ion (HO)2PO2- and the related ions, diphosphate, D-ribose 5-phosphate, adenosine 5-phosphate, and adenosine 5-diphosphate. The hydration energies of singly and multiply charged anions such as anions derived from nucleic acids and adenosine phosphates are of interest for more than one reason. (a) The availability of such data will aid in the theoretical modeling of these ionic compounds in aqueous solution. Such modeling efforts13 applied to multiply charged peptides, proteins, and nucleic acids represent an important method in deducing X
Abstract published in AdVance ACS Abstracts, January 15, 1996.
0022-3654/96/20100-2443$12.00/0
structural differences between crystal diffraction X-ray deduced structures and structures in a native aqueous environment. (b) Another important use of the data is connected to recent mass spectrometric work for the determination of molecular weights and sequencing of singly and multiply charged oligonucleotides.14 The presence of single and multiple charges is expected to lead to internal stabilization of the gas-phase ions by cyclization due to intramolecular H-bond formation. This situation is analogous to the more widely recognized and studied internal H-bond formation in protonated peptides.11,15,16 The formation of strong intramolecular H bonds in gas-phase ions such as the protonated R,ω-diamines NH3+(CH2)kNH2 was first convincingly demonstrated17 by studies of proton-transfer equilibria involving the R,ω-diamines. Proton transfer to these compounds led to a much higher proton affinity than expected for protonation of the single amino group, indicating dicoordination of the proton by the two amino groups. Furthermore, the entropy change on protonation was negative and the value was consistent with the loss of freedom due to the cyclization. Proton-transfer equilibria determined at different temperatures are the best method of establishing the occurrence of internal H bonding since these measurements provide the enthalpy and entropy change.17,18 Equilibria determinations at a single temperature are also very useful indicators of the presence of cyclization. In that case internal dicoordination and cyclization are inferred from an exoergicity than is larger than expected for a single basic functional group. Proton-transfer equilibria cannot be applied to charged polypeptides or polynucleotides because the neutral precursors of these ions are of very low volatility and cannot be introduced into the ion reaction chamber at sufficient partial pressures. In such cases deprotonation of the protonated peptides by volatile bases of known gas-phase basicities, the bracketing methods, has been used.15 On this basis and theoretical calculations, Cassady and co-workers15 concluded that internal H bonding via cyclization is a prominent factor affecting the structure of protonated peptides. The effect of hydrogen bonding on the structure of polyprotonated peptides and proteins has been subject to recent systematic studies by Williams and coworkers.16 Determinations of hydration energies of ions represent an alternate way of establishing the presence of internal stabilization © 1996 American Chemical Society
2444 J. Phys. Chem., Vol. 100, No. 6, 1996
Blades et al.
by H bond formation. For example, the cyclized protonated R,ω diamines have much lower hydration energies.10-12,18 The same is true13 for deprotonated cyclized R,ω diacids such as HO2C(CH2)kCO2-. Furthermore, when the ring is strained, as is the case for short CH2 chains, and the internal H bond is weak, relatively higher hydration energies are observed.10-13 In other words, the magnitudes of the hydration energies can be a very good probe for the presence and strength of internal H bonding.10-12 The presence of internal H bonding in the sugar phosphates and nucleotide anions are the main subject of the present work. Experimental Section The hydration equilibria (eq 1) were determined with an apparatus which has been described in detail.11,12 Therefore only the main features will be given here. A solution containing ∼10-4 mol/L of the sodium phosphate in methanol solvent was infused through the electrospray capillary, at ∼1 µL/min flow rate, by means of a motor-driven syringe. The electrospray capillary which is at high negative voltage (-4.5 kV) emits a spray of droplets which are negatively charged due to an excess of anions over cations in the droplets. Droplet evaporation in the atmosphere leads to gas-phase negative ions. Part of the spray is sucked in by a second capillary (-40 V) leading to a chamber (-40 V) maintained at 10 Torr. The gas jet and ions entering the low-pressure chamber pass through an electric field which deflects the ions into the reaction chamber (-17 V). A gas mixture of 10 Torr bath gas N2 and a known partial pressure of water vapor flows through the reaction chamber. Ion equilibria establish in this chamber; for discussion of equilibrium conditions, see previous publications.12 The relative ion concentrations are determined by letting a fraction of the gas and ions escape, through an orifice (-3 V), into a vacuum chamber which houses a triple quadrupole mass spectrometer. The reaction chamber, in which the ion equilibria occur, was not provided with heaters or cooling channels. This ion source was operated at a single temperature, which is close to the airconditioned room temperature. The temperature of the chamber (293 K), determined with a thermocouple,12 was somewhat lower than the room temperature due to some radiative cooling of the chamber by the cryopump surfaces of the vacuum system.12 The following salts were used: Na2HPO4 produced (HO)2PO2-, Na4P2O7(diphosphate) produced mainly the dianion (HO)O2POPO2(OH)22-. Na2-D-ribose 5-phosphate produced the singly charged anion. Na-adenosine monophosphate led to the singly charged ion, while the disodium salt of adenosine diphosphate gave a good yield of the doubly charged ion. Results and Discussion (a) Results. The equilibrium constants, Kn-1,n, for the hydration of the phosphate anions (eq 2) were obtained from mass
AZ-(H2O)n-1 + H2O ) AZ-(H2O)n
(2)
spectrometric determinations of the ion intensity ratio at a known partial pressure of H2O, PH2O.
Kn-1,n )
In In-1pH2O
(3)
The ion intensity ratio, In/In-1, is assumed to be equal to the ion equilibrium concentration ratio, [AZ-(H2O)n]/[AZ-(H2O)n-1]. For
Figure 1. Plot of ion intensity ratio In/In-1 of ion hydrates versus partial pressure of water, PH2O. According to eq 3, plot should be linear and go through origin. The slope provides the equilibrium constant Kn-1,n. Plots for I2/I1 of the (HO)2PO2- ion and I1/I0 of the D-ribose 5-phosphate singly charged ion are shown.
details see Experimental Section and previous work.10-12 Representative plots of In/In-1 versus PH2O are shown in Figure 1. According to eq 3, the slope of the plots should be constant and equal to Kn-1,n. Good linear plots are obtained in Figure 1 which also go through the origin as required by eq 3. The plots shown exhibit a slight deviation from linearity at higher pressure. The reason for this deviation is not known. Often in experiments with other systems better linearity is observed, see for example Figure 7 in ref 11. Considering that the water vapor pressure is increased by the very large factor of 25 and a chemical equilibrium is involved which is determined mass spectrometrically, the observed deviations in Figure 1 can be considered to be small relative to previous ion-molecule equilibria determinations obtained with other apparatus and reported in the literature.2 For the evaluation of the equilibrium constants we have selected the slope observed at low pressure, as indicated by the straight lines drawn in Figure 1. The two plots for K1,2 of (HO)2PO2- and K0,1 of D-ribose 5-phosphate were selected not only because the two equilibrium constants are of similar magnitude and fit well in the same plot but also because these two determinations are of special interest; see section b. The equilibrium constants and free energy changes obtained with the equation
-∆G°n-1,n ) RT ln Kn-1,n
(4)
are, at 293 K
(HO)2PO2K1,2 ) 35 700 atm-1 -∆G°1,2 ) 6.1 kcal/mol
D-ribose
5-phosphate
K0,1 ) 22 000 atm-1 -∆G°0,1 ) 5.8 kcal/mol
Hydration free energies, ∆G°n-1,n, for the phosphates obtained from plots like those shown in Figure 1, are summarized in Table 1. Also given in Table 1 are hydration free energies of some sulfate and carboxylate anions from our previous negative ion determinations.12 These data will be useful for the discussion of the present results. (b) Hydration Energies of Some Phosphate Anions Including D-Ribose 5-Phosphate, and Adenosine 5-Monophosphate. Evidence for Intramolecular H bonding. The two plots leading to the equilibrium constants, K1,2, of (HO)2PO2and K0,1 of ribose phosphate (Figure 1) illustrate at a glance a
Free Energies of Hydration
J. Phys. Chem., Vol. 100, No. 6, 1996 2445
TABLE 1: Hydration Free Energies in Gas Phasea Singly Charged Anions (-∆G°n-1,n) ions (n - 1,n) (0,1) (1,2) H2PO28.4 6.4 (HO)HPO27.9 6.6 (HO)2PO27.6 6.1 D-ribose 5-phosphate 5.8 4.5 adenosine 5′-phosphate 5.4b HOSO3 5.9 4.6 C2H5OSO35.9 4.6 C2H5CO29.3 6.6 HOCO(CH2)5CO25.5 Doubly Charged Anions (-∆G°n-1,n) (HO)O2POPO2(OH)2(2,3) 8.9 (3,4) 7.9 (adenosine 5′-diphosphate) (0,1) 9.3 (1,2) 7.9 (4,5) 5.6 (5,6) 5.1 SO42(5,6) 8.5 (6,7) 7.5 O3SOOSO32(1,2) 8.5 (2,3) 7.5
(2,3) 4.8 5.2 4.8
5.1
(4,5) 6.7 (2,3) 6.8 (6,7) 4.6 (7,8) 6.7 (3,4) 6.7
a Free energy changes at 1 atm and 293 K for reaction AZ-(H2O)n-1 + H2O ) AZ-(H2O)n-1, where Z ) 1, 2, in kcal/mol. The table includes also values for sulfates and carboxylates, obtained in previous work,11 which provide useful comparisons. An experimental error for the free energy determination is estimated to be less than (0.2 kcal/mol. b Preliminary measurements of the ∆G° 0,1 for adenosine 3′-phosphate provide a value which is 0.1 kcal/mol lower than that for the adenosine 5′-phosphate.
point that will be central to the discussion. The plots show that for the same conditions (HO)2PO2- adds on one more water molecule and is thus more strongly hydrated than the ribose phosphate. Previous experimental work has shown that a weaker hydration for a given anion is observed when (a) the negative charge is delocalized over more, equivalent oxygen atoms, (b) when electron withdrawing or more polarizable substituents are present, and (c) when there is stabilization of the anion by an intramolecular hydrogen bond.12 The phosphate and ribosephosphate anions have the same number (2) of equivalent oxygens over which charge is delocalized and therefore, the weaker hydration of the ribose phosphate should be due to the factors described in (b) and/or (c). The magnitude of the electron withdrawing effect of the ribose relative to OH can be estimated on the basis of data shown in Table 1. The decrease of the solvation energies (-∆G0,1) in the series H2PO2- (8.4), (HO)HPO2- (7.9), (HO)2PO2- (7.6), where there are two equivalent oxygens over which the charge is delocalized, can be attributed primarily to the electron-withdrawing effect of the hydroxy groups. The electron-withdrawing effect of a hydroxy group is similar to that of an alkoxy group as illustrated by the (-∆G°0,1) values for HOSO3- (5.9) and C2H5OSO3- (5.9) given in Table 1. The electron-withdrawing effect of ribose should be similar to that of an alkoxy substituent because the electron-withdrawing OH groups on the ribose are relatively remote to have any significant effect. The polarizability of ribose will be substantially higher than that of the ethyl group, however, the polarizability of groups at relatively remote locations has only a small effect on the strength of the hydration. This is illustrated by the -∆G°0,1 values for CH3SO3- (6.9) and n-C7H15SO3- (6.4) determined in previous work.12 On the basis of the above considerations, we conclude that the large difference between the -∆G°0,1 values of the phosphate (7.6) and ribose phosphate (5.4) is mostly due to the formation of an intramolecular hydrogen bond in the ribose phosphate. A comparison can be made with the hydration energies (-∆G°0,1) of the propionate anion C2H5CO2- (9.3) and the R,ωdicarboxylate anion, HOCO(CH2)5CO2- (5.5). The dicarboxy-
late is known to be stabilized by an internal H bond involving the two terminal carboxy groups.11,12 The hydration energy difference between the uncyclized, no internal H bond, C2H5CO2and the cyclized HOCO(CH2)5CO2- is somewhat larger than the observed difference for the phosphates, (HO)2PO2-, and ribose phosphate. This observation is consistent with a smaller stabilizing effect of the internal H bond for the phosphates. The result is expected. The strength of the H bond formed depends on the acidity of the hydrogen donor and the basicity of the acceptor.19 For the ribose phosphate, the H donor is a hydroxy group which is less acidic than the OH of a carboxylic group. Also, the base, i.e., the hydrogen acceptor, is weaker because the anionic phosphate group is more stabilized, i.e., a weaker base, than the anionic carboxylic group. An examination with space filling structures indicates that the spatial requirements for an intramolecular H bond in the D-ribose 5-phosphate are best met by the hydroxy group in position 3 of the ribose, as shown in the schematic structure I.
The hydration energy of adenosine 5′-monophosphate -∆G°0,1 ) 5.4 kcal/mol (Table 1) is also low compared to that of (HO)2PO2- and an analogous intramolecular H bond must be present also in this ion; see structure II. The hydration energy of the adenosine 3′-monophospshate was found to be very close to that for the 5′-monophosphate (Table 1). This indicates that also the 3′-monophosphate forms an intramolecular H bond. An eight-membered ring is formed by the 5′-phosphate, while the 3′-phosphate forms a seven-membered ring. One expects similar stabilities for these two H-bonded systems and this is supported by the closeness of the hydration values. Deprotonated nucleotides subjected to collision-induced decomposition lead to a number of product ion types.14 An important group are the deprotonated bases such as the adenilyl negative ion when the base is adenine. The mechanism for the formation of these ions has not been established. A mechanism that has been suggested14b,d involves an intramolecular nucleophilic substitution in which a phosphate oxygen displaces the base. Phillips and McCloskey14b have proposed that this reaction is facilitated by the formation of a hydrogen bond between the other phosphate oxygen and a hydroxide group of the sugar which carries the base; see Scheme 7 in ref 14b. The hydrogen-bonded structure involves the same interaction as proposed here for the adenosine-3′-monophosphate. (c) Hydration Results for Doubly Charged Diphosphates. Hydration data for the doubly charged diphosphates are given in Table 1. Previous work12 has shown that doubly charged ions have much higher hydration energies than singly charged ions. As was the case for singly charged ions, the hydration energies of doubly charged ions decrease as the charge is delocalized. A very effective charge delocalization occurs when the charge is separated onto two singly charged groups and the distance between the groups is increased.12 An example of this effect is seen in the data12 for the sulfate, SO42-, and the persulfate O3SOOSO32-, anions quoted in Table 1. The hydration energies become approximately the same only after the sulfate ion has acquired four water molecules.
2446 J. Phys. Chem., Vol. 100, No. 6, 1996 The hydration of the diphosphate (HO)O2POPO2(OH)2- is found to be stronger than that for the persulfate O3SOOSO32-. This should be due to the shorter distance between the two charged groups in the diphosphate and the lesser charge delocalization per group present in the diphosphate. Thus per group, there are only two charged oxygens in the phosphate versus three in the sulfate. This comparison also indicates that there is no internal stabilization by cyclization where one OH group of the one phosphate group interacts with negative oxygens of the other phosphate group. Such an interaction, if favorable, would have brought the hydration energies of the diphosphate and persulfate, which has no OH groups, closer. The cyclization probably does not occur because of increased Coulombic repulsion between the two charged groups in the cyclized diphosphate. The hydration of the adenosine 5′-diphosphate is observed to be much weaker than that of the diphosphate (Table 1). The hydration energies become approximately the same only after the diphosphate has acquired two more water molecules. The much weaker hydration of the adenosine diphosphate must be due to formation of an intramolecular hydrogen bond between an anionic oxygen on the terminal phosphate and a hydroxy group of the ribose. Space-filling models again indicate that the most favorable interaction occurs with the 3′-hydroxy group on the ribose, although confirmation of this structure on the basis of molecular modeling or, even better, theoretical calculations would be required. We plan to extend the hydration measurements to adenosine triphosphate and to oligomeric nucleic acids. A newly constructed ion source which allows equilibria determinations above and below room temperature will allow us to extend the range of sequential hydration steps that we can determine and also allow, through van’t Hoff plots, to obtain enthalpy, ∆H°n-1,n, and entropy, ∆S°n-1,n values. Conclusions Hydration energies of some singly and doubly charged phosphates in the gas phase have been determined using ions produced by electrospray. These data will be useful in attempts to model the behavior of these ions in aqueous environments, in particular, those of biological significance. The hydration energies also provide information regarding the presence of stabilization by intramolecular hydrogen bonding in the unhydrated ions. Cyclization by intramolecular hydrogen bonding can have significant structural effects, particularly for
Blades et al. multiply charged ions such as oligonucleic acids. The cyclized structures may be of importance in the condensed phase when a poorly solvating hydrophobic environment is present. The presence of these structures in the gas phase may be also of importance in mass spectrometric bioanalytical work where sequencing of nucleic acids is obtained by collision induced decomposition of the multiply charged nucleic acids. The situation is somewhat analogous to the more widely explored effects of cyclization due to intramolecular hydrogen bonding occurring in protonated and polyprotonated peptides.11,15,16 References and Notes (1) Hogg, A. M.; Haynes, R. N.; Kebarle, P. J. Am. Chem. Soc. 1966, 88, 28. (2) Kebarle, P. Annu. ReV. Phys. Chem. 1977, 28, 455. (3) Keesee, R. G.; Castleman, A. W. J. Phys. Chem. Ref. Data 1986, 15, 1011. (4) Yamashita, M.; Fenn, J. B. J. Phys. Chem. 1984, 88, 4451. Fenn, J. B.; Mann, M.; Meng, C. K.; Wong, S. F.; Whitehouse, C. M. Science 1985, 246, 64. (5) Kebarle, P.; Tang. L. Anal. Chem. 1993, 65, 272A. (6) Blades, A. T.; Jayaweera, P.; Ikonomou, M. G.; Kebarle, P. J. Chem. Phys. 1990, 92, 2900. Blades, A. T.; Jayaweera, P.; Ikonomou, M. G.; Kebarle, P. Int. J. Mass Spectrom. Ion Processes 1990, 102, 251. (7) Bruins, A. P.; Covey, T. R.; Henion, J. D. Anal. Chem. 1987, 59, 2642. (8) Blades, A. T.; Kebarle, P. J. Am. Chem. Soc. 1994, 116, 10761. (9) Limbach, P. A.; Crain, P. F.; McCloskey, J. A. J. Am. Soc. Mass Spectrom. 1995, 6, 27. Smith, R. D.; Lou, J. A.; Edonds, C. G.; Barinaga, C. J.; Udseth, H. R. Anal. Chem. 1990, 62, 882. (10) Klassen, J. S.; Blades, A. T.; Kebarle, P. J. Am. Chem. Soc. 1994, 116, 12075. (11) Klassen, J. S.; Blades, A. T.; Kebarle, P. J. Phys. Chem. 1995, 99, 15509. (12) Blades, A. T.; Klassen, J. S.; Kebarle, P. J. Am. Chem. Soc. 1995, 117, 10563. (13) See for example: Gao, J.; Keczera, K.; Tidor, B.; Karplus, M. Science 1989, 244, 1069. (14) (a) Pomeranz, S. C.; Kowalak, J. A.; McCloskey, J. A. J. Am. Soc. Mass Spectrom. 1993, 4, 204. (b) Phillips, D. R.; McCloskey, J. A. Int. J. Mass Spectrom. Ion Processes 1993, 128, 61. (c) McLucky, S. A.; HabibiGondazzi, S. J. Am. Chem. Soc. 1993, 115, 12085. (d) Rodgers, M. T.; Campbell, S.; Marzluff, E. M.; Beauchamp, J. L. Int. J. Mass Spectrum. Ion Processes 1994, 137, 121. (15) Zhang, K.; Zimmerman, D. M.; Chung-Phillips, A.; Cassady, C. J. J. Am. Chem. Soc. 1993, 115, 10812. (16) Gross, D. S.; Williams, E. R. J. Am. Chem. Soc. 1995, 117, 883. Gross, D. S.; Rodriguez-Cruz, S. E.; Bock, S.; Williams, E. R. J. Phys. Chem. 1995, 99, 4034. (17) Yamdagni, R.; Kebarle, P. J. Am. Chem. Soc. 1973, 95, 3504. (18) Meot-Ner (Mautner), M.; Hamlet, P.; Hunter, E. P.; Field, F. H. J. Am. Chem. Soc. 1980, 102, 6393. (19) Paul, J. C.; Kebarle, P. Can. J. Chem. 1990, 68, 2070.
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