Deducing Proton-Bound Heterodimer Association Energies from Shifts

Subscriber access provided by UNIV OF TEXAS DALLAS

A: New Tools and Methods in Experiment and Theory

Deducing Proton-Bound Heterodimer Association Energies from Shifts in Ion Mobility Arrival Time Distributions Pearl Kwantwi-Barima, Christopher J. Hogan, and Brian H Clowers J. Phys. Chem. A, Just Accepted Manuscript • DOI: 10.1021/acs.jpca.8b11183 • Publication Date (Web): 07 Mar 2019 Downloaded from http://pubs.acs.org on March 16, 2019

Just Accepted “Just Accepted” manuscripts have been peer-reviewed and accepted for publication. They are posted online prior to technical editing, formatting for publication and author proofing. The American Chemical Society provides “Just Accepted” as a service to the research community to expedite the dissemination of scientific material as soon as possible after acceptance. “Just Accepted” manuscripts appear in full in PDF format accompanied by an HTML abstract. “Just Accepted” manuscripts have been fully peer reviewed, but should not be considered the official version of record. They are citable by the Digital Object Identifier (DOI®). “Just Accepted” is an optional service offered to authors. Therefore, the “Just Accepted” Web site may not include all articles that will be published in the journal. After a manuscript is technically edited and formatted, it will be removed from the “Just Accepted” Web site and published as an ASAP article. Note that technical editing may introduce minor changes to the manuscript text and/or graphics which could affect content, and all legal disclaimers and ethical guidelines that apply to the journal pertain. ACS cannot be held responsible for errors or consequences arising from the use of information contained in these “Just Accepted” manuscripts.

is published by the American Chemical Society. 1155 Sixteenth Street N.W., Washington, DC 20036 Published by American Chemical Society. Copyright © American Chemical Society. However, no copyright claim is made to original U.S. Government works, or works produced by employees of any Commonwealth realm Crown government in the course of their duties.

Page 1 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

The Journal of Physical Chemistry

Deducing Proton-Bound Heterodimer Association Energies from Shifts in Ion Mobility Arrival Time Distributions Pearl Kwantwi-Barima,1 Christopher J. Hogan Jr.,2 Brian H. Clowers1,* 1

Department of Chemistry, Washington State University, Pullman, WA 99164

2

Department of Mechanical Engineering, University Minnesota, Minneapolis, MN

* Corresponding Author: [email protected] Phone: 509-335-4300 ABSTRACT Through vapor modification of the counter-current drift gas in an atmospheric pressure drift tube ion mobility spectrometer (IMS), we demonstrate measurement of gas-phase association enthalpies and entropies for select proton bound heterodimers formed from a phosphonic acid with 2-propanol. Previous efforts to determine gas-phase association thermodynamic properties have relied largely upon lower pressure systems and inference of the relative concentrations of m/z isolated species. In contrast, the drift tube IMS based approach developed and applied in this study leverages the explicit gasphase equilibrium that is established within an ion mobility drift cell.

The inferred

enthalpies and entropies of association are based solely upon monitoring a shift in the arrival time of an ion at different temperatures (and not on the signal intensity or on external instrument drift time calibration). We specifically report the gas-phase Gibbs free energy, enthalpy, and entropy changes for the association of 2-propanol with protonated methyl, ethyl, and propyl phosphonic acid ions (MPA, EPA, PPA) across the 100-175 oC temperature range. For all of these proton bound heterodimers, the standard enthalpies and entropies of 2-propanol association were negative and positive, respectively. These data indicate that proton bound heterodimer formation is both enthalpically and entropically favorable, though we find that the magnitude of the standard enthalpy change for vapor association is small (near 1 kcal/mol for all examined heterodimers). Though many prior results (largely obtained with high pressure massspectrometry) for other proton-bound organic heterodimer complexes show larger enthalpic favorability and an entropic barrier, our results qualitatively conform to the bulk

ACS Paragon Plus Environment

1

The Journal of Physical Chemistry 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 2 of 32

Kelvin-Thomson-Raoult (KTR) model, which is commonly utilized in describing ioninduced nucleation of a vapor onto a soluble, nanometer scale ion. The KTR model suggests that heterodimer formation due to vapor binding to an ion should be slightly enthalpically favored (due to a larger Thomson effect than the Kelvin effect) and entropically favored because of ion solvation (Raoult’s effect). The method presented in this study can be applied to any static-field ion mobility spectrometer and to a wide variety of heterodimers. Due to the ease of implementation and broad applicability, this approach may find consistent use in determining the thermodynamic properties of weakly bound gas phase heterodimer complexes which are difficult to probe via alternative techniques. Moreover, this renewed implementation of the IMS experiment is directly compatible with soft ionization sources which will enable the characterization of vapor modifier induced mobility shifts experiments for larger molecular complexes.

2 ACS Paragon Plus Environment

Page 3 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


INTRODUCTION The design of ion mobility spectrometry (IMS) systems and interpretation of IMS measurements is based upon a description of the interactions between ions and an inert buffer gas in the presence of a weak electric field .1,2,3,4 In a drift tube IMS system, the electric field drives ion motion, which is resisted by the buffer gas, and the mobility is the proportionality coefficient between the ion’s steady-state velocity and the electric field strength. A variety of methods exist to exploit mobility for separation of ions, and modern manufacturing techniques enable rather inexpensive IMS design and production.5 For these reasons, atmospheric pressure IMS finds extensive use in field settings for the identification and separation of explosives,6 chemical warfare agents,7,8 and illicit narcotics.9 When coupled with soft ionization sources, ion mobility coupled with mass spectrometry also enables structural characterization of peptides, proteins, and other biopolymers in the gas phase.10,11 Regardless of the application, enhancing the separation capacity and selectivity of IMS, as well as increasing the amount of information which can be extracted from IMS measurements, would be of considerable benefit. There are, however, challenges in separation capacity enhancement and improving chemical characterization in IMS. A central issue is that an ion’s mobility is most strongly dependent upon its overall structure,12 and while it is also affected by the drift gas molecular structure and the ion charge state, it is only weakly dependent upon the true chemical nature of the ion (i.e. the chemical bonds and functional groups constituting the ion). A possible method to enhance chemical selectivity and separation capacity in IMS is drift gas composition modulation. In an analogous fashion to ion-pairing agents in chromatography

13,14

, the

introduction of condensable drift gas modifiers, even at low concentrations, alters the traditional separation factors for homogeneous drift gas experiments. Such modifiers can not only collide with and transfer momentum to ions (as non-condensable drift gas molecules do), but can also transiently bind to analyte ions forming heterodimers or higher order clusters, and in doing so shifting apparent mobilities. As the extent of binding



Page 4 of 32

is strongly dependent upon on the chemical interactions of the ion and modifier, 15,16 the change in an ion’s mobility as function of modifier concentration is directly proportional to concentration of the associating vapor and the favorability of this interaction. Although investigations of vapor modifier induced mobility shifts have been previously reported,17-18 only selected efforts have been directed towards quantifying the extent of mobility shift and establishing and a core predictive capacity. Works by Rawat et al.19 and Oberreit et al.20 established a framework to quantitatively account for vapormediated mobility shifts occurring at ambient pressure, based upon the assumption that ions are in equilibrium with the surrounding drift gas in an IMS system, and that drift gas modifier molecules sorb and desorb from the ion as it traverses an ion mobility spectrometer (hence the number of bound vapor molecules is not fixed). Recently, work by Kwantwi-Barima et al.21 refined the Oberreit et al.20 model to specifically account for single modifier molecule-induced mobility shifts at atmospheric pressure. This particular model linked the observed mobility shift in the presence of a drift gas modifier to the equilibrium binding coefficient for a single drift gas modifier molecule binding transiently to the ion, the collision cross section of the bare analyte ion, and the collision cross section of the analyte ion-drift gas modifier complex. Through its direct link to the equilibrium binding coefficient, the aforementioned model enables direct inference of the Gibbs free energy change for the formation of heterodimer ions (original ion and vapor molecule) in the gas phase. Though not yet explored, this model additionally suggests that by making measurements of mobility shifts for an ion at variable drift gas modifier concentrations and variable system temperature, it is possible to experimentally infer the changes in enthalpy and entropy of association for heterodimer complexes. This would present a new and potentially widely applicable approach to characterize weakly bound heterodimers and ion-molecule clustering in the gas phase. To date, thermodynamic properties of cluster ions and heterodimers have been determined primarily through the use of high pressure mass spectrometric


Page 5 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


techniques,22 ion cyclotron resonance (ICR),23 and a combination of equilibrium thermochemical measurements and ion mobility techniques.24,25 While such efforts have enabled determination of the enthalpies and entropies of association for wide variety of mono- and polyatomic neutrals onto atomic ions and inorganic ions, and water onto organic ions,26

27 28 29 30 31 31 32 33 34,35 36,

these methods require relatively specialized

equipment. Moreover, heterodimer formation from organic polyatomic neutrals with polyatomic organic ions has largely been limited to ions generated from high vapor pressure species,33 as opposed to those generated from more recently developed soft ionization sources for analytes derived from the condensed phase. The ability to assess the enthalpies and entropies of association on more common IM-MS instruments (i.e. on linear IM-MS instruments with minimal change to instrument operating procedures) for a wide variety of analytes would aid in better understanding ion dynamics in the vapor phase and would aid in the “design” of IM-MS based separation schemes with improved chemical selectivity. In this work we carry out mobility shift experiments from 100 oC to 175 oC, and for these conditions we assess the gas-phase association energies for heterodimer formation from protonated ions of methyl phosphonic acid (MPA), ethyl phosphonic acid (EPA), propyl phosphonic acid (PPA) with 2-propanol (the drift gas modifier vapor). Distinct from prior efforts using high pressure mass spectrometry and selected ion flow tube systems,37,38 the approach detailed in the current work focuses on relative peak location to deduce thermodynamic values, rather than peak areas and intensities, and appears to be uniquely well-suited to examine the thermodynamic properties of weakly-bound, transiently-formed, proton-bound heterodimers from a wide variety of precursor ions and neutrals.

EXPERIMENTAL Chemicals and Reagents. Vapor-mediated shifts in mobility were evaluated for protonated ions of methyl phosphonic acid (MPA), ethyl phosphonic acid (EPA), and propyl phosphonic acid (PPA). Isopropyl alcohol (2-propanol) was used as the drift gas



Page 6 of 32

modifier. In addition to the analytes and the gas-phase modifier listed above, HPLC grade methanol and 0.1% formic acid (ACS reagent grade, >= 97 or 98% purity) were purchased from Sigma-Aldrich Chemical Co. (Milwaukee, WI, U.S.A.). All analytes (MPA, EPA, PPA) were formulated at a concentration of 10 µM in HPLC grade methanol with 0.1% formic acid. Ion Mobility-Mass Spectrometry. Mobility and mass spectral information were obtained using an atmospheric pressure, dual-gate drift tube ion mobility spectrometer (ExcellIMS, MA3100) coupled to a linear ion trap mass spectrometer, LTQ (Thermo Fisher Scientific, Thousand Oaks, CA), which was also used in our previous study.21 The target analytes were infused through a 75 µm glass capillary at a flow rate of 2 µL/min. The electrospray emitter was held at a potential of 2400 V (relative to first IMS electrode). Ions were gated using a Bradbury-Nielsen ion gate with an interwire potential of +/- 50 V. The potential of the drift cell was held at 8000 V (unless stated otherwise) and operated under atmospheric pressure (∽ 690 Torr in Pullman WA). A clean nitrogen flow of 1.5 L/min was mixed with vapor modifier laden drift gas at 0.5 L/min; these flows entered the drift tube, where they were held at a temperature of 100 - 175 +/- 0.2 oC. Using a custom pulled glass capillary as an electrospray source, ion current was measured both with a Faraday plate and LTQ mass analyzer. By using the custom ExcellIMS electronics assisting the MA3100 unit, an electric field of ∽419 V cm-1 was utilized throughout the experiment. Mass-selected mobility spectra were obtained by frequency encoding the mobility data using the approach detailed by Morrison et al.39 For the frequency modulated IMS experiments (Fourier Transform-IMS (FT-IMS)) the synchronized gating frequency was swept from 5 Hz - 10.5 kHz over the course of 8 minutes. Drift Gas Modifier Introduction. 2-propanol was introduced as the drift gas modifier into the primary drift gas inlet through a glass capillary coupled to a temperature-controlled GC injection liner at a variable flow rate from 5 µL/hr to 160 µL/hr using a syringe pump. To avoid the potential for temperature gradients, the temperature of the drift gas mixing


Page 7 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


chamber was kept at the same temperature as the ion mobility drift tube. In addition to the above conditions, the gas phase saturation ratio range for 2-propanol at the 𝑃

aforementioned flow rate corresponds to a saturation ratio of 𝑃 (𝑇)= 0-2.43 x 10-5, 0𝑠𝑎𝑡 4.71 x 10-5, 0-9.97 x 10-5, 0-2.34 x 10-4 for 175 oC, 150 oC, 125 oC and 100 oC respectively. The saturation ratio (S) is defined as the ratio of actual experimental vapor pressure at the measurement temperature divided by the saturated vapor pressure of the vapor modifier at the measurement temperature (T). The actual experimental vapor pressure of the modifier at the measurement temperature was calculated from the experimental vapor concentration using the ideal gas law. In order to demonstrate reproducibility of the identified shifts, experiments were conducted in triplicate in a sequential order for all drift gas modifier conditions.

COMPLEX GAS-PHASE EQUILIBRIA SHIFT MODEL Our prior work provides the link between the change in drift time (or change in mobility), the collision cross section of a bare ion, the collision cross section of an ion-molecule complex, and the Gibbs free energy change associated with binding.

21

In the present

set of experiments, we extend that effort by evaluating the mobility shift model across different temperatures and substituting the saturation ratio for the normalized vapor concentration of the drift gas-phase modifier. The latter quantity is the actual experimental vapor pressure of the modifier at the measurement temperature divided by the saturated vapor pressure of the vapor modifier at a reference temperature (298 K). Using the saturation ratio in our previous publication was certainly an appropriate choice for a reference concentration in a fixed temperature experiment. However, in order to compare our measured thermodynamic properties to literature values (even qualitatively),33 the vapor pressure of the modifier at different measurement temperatures must be compared to a reference standard state at a defined constant temperature, as has been done in high pressure mass spectrometry studies.29,33 This



Page 8 of 32

leads to a modification of the model reported by Kwantwi-Barima et al.21 as follows. Equation 1 relates the arrival time for an ion, t, in the presence of a vapor modifier at saturation ratio through the relationship: 𝑡 𝑡0

(

1 + 𝑆𝑒𝑥𝑝 ―

=

𝛺0𝜇1/2 1

1 + 𝑆𝛺

∆𝐺1,𝑆

(

𝑘𝑇

𝑒𝑥𝑝 ―

1𝜇1/2 0

)

∆𝐺1,𝑆 𝑘𝑇

(1)

)

where, t0, is the initial arrival time measured in the absence of the vapor modifier, Ω0 is the collision cross section of a bare ion, Ω1 is the collision cross section of the ion-vapor molecule complex (here, proton bound heterodimers), µ0 is the reduced mass of the ionbath gas system, µ1 is the reduced mass of the ion-neutral complex-bath gas system, and finally, ∆𝐺𝑖,𝑆 is the Gibbs free energy change associated with vapor uptake of the “ith” vapor molecule under saturated conditions at the measurement temperature T. 𝑆 is given by the equation: 𝑃

(2)

𝑆 = 𝑃𝑠𝑎𝑡(𝑇) where 𝑃𝑠𝑎𝑡(𝑇) is the saturated vapor pressure of the modifier at the measurement temperature T and P is the experimental vapor pressure calculated from the experimental vapor concentration. From the Clausius-Clapeyron equation:

(

𝑃𝑠𝑎𝑡(𝑇) = 𝑃𝑟𝑒𝑓(𝑇𝑟𝑒𝑓)𝑒𝑥𝑝 ―

𝐻𝑣𝑎𝑝 1 𝑘

[

1

])

𝑇 ― 𝑇𝑟𝑒𝑓

(3)

𝑃𝑟𝑒𝑓(𝑇𝑟𝑒𝑓) is the saturated vapor pressure of the modifier at a specifically defined reference temperature 𝑇𝑟𝑒𝑓and 𝐻𝑣𝑎𝑝 is the enthalpy of vaporization for the vapor modifier (10.86 kcal mol-1 for 2-propanol). By substituting equations (2) and (3) into equation (1), the drift time ratio can be expressed as follows:


Page 9 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


1+𝑃

𝑡 𝑡0

= 1+𝑃

𝑃

(

)𝑒𝑥𝑝 ―

𝑟𝑒𝑓(𝑇𝑟𝑒𝑓

𝛺0𝜇1/2 1

𝑃


(―

𝛺 1/2𝑒𝑥𝑝 𝑟𝑒𝑓(𝑇𝑟𝑒𝑓) 1𝜇0

+

𝐻𝑣𝑎𝑝 1


𝑘

+

[

𝑇

―𝑇

𝐻𝑣𝑎𝑝 1 𝑘

[

𝑇

1

])

𝑟𝑒𝑓

―𝑇

1

])

(4)

𝑟𝑒𝑓

Consequently, the Gibbs free energy of association based on the reference state the relationship then becomes:

𝑡 𝑡0

1+𝑃

=

1+𝑃

𝑃

(

)𝑒𝑥𝑝 ―

𝑟𝑒𝑓(𝑇𝑟𝑒𝑓

𝑃

𝛺0𝜇1/2 1

∆𝐺1,𝑟𝑒𝑓 𝑘𝑇

(―

𝛺 1/2𝑒𝑥𝑝 𝑟𝑒𝑓(𝑇𝑟𝑒𝑓) 1𝜇0

)

∆𝐺1,𝑟𝑒𝑓 𝑘𝑇

(5)

)

Comparing Equations (4) and (5) a linear relationship is produced: ∆𝐺1,𝑟𝑒𝑓 𝑘𝑇

=


―

𝐻𝑣𝑎𝑝 1 𝑘 𝑇

[

1

]

(6)

― 𝑇𝑟𝑒𝑓

Therefore, at any given temperature, the Gibbs free energy inferred from measurement based on a consistent reference state (∆𝐺1,𝑟𝑒𝑓) is linked to the Gibbs free energy at a variable saturation state via the equation:

(

𝑇

)

(7)

∆𝐺1,𝑟𝑒𝑓 = ∆𝐺1,𝑆 ― 𝐻𝑣𝑎𝑝 1 ― 𝑇𝑟𝑒𝑓

Correspondingly, we can define changes in enthalpy and entropy under saturated conditions, and with respect to a constant temperature reference state: ∆𝐺1,𝑆 = ∆𝐻1,𝑆 ―𝑇∆𝑆1,𝑆

(8)

∆𝐺1,𝑟𝑒𝑓 = ∆𝐻1,𝑟𝑒𝑓 ―𝑇∆𝑆1,𝑟𝑒𝑓

(9)

Grouping terms yields: (10)

∆𝐻1,𝑟𝑒𝑓 = ∆𝐻1,𝑆 ― 𝐻𝑣𝑎𝑝



∆𝑆1,𝑟𝑒𝑓 = ∆𝑆1,𝑆 ―

𝐻𝑣𝑎𝑝

Page 10 of 32

(11)

𝑇𝑟𝑒𝑓

With these equations, measurements are then analyzed as follows. First, we group measurements by temperature, and fit Equation (1) to plots of the drift time ratio (t/ti) as 𝑃

a function of saturation ratio (𝑃 (𝑇))to yield ∆𝐺1,𝑆 at each examined temperature, for 𝑠𝑎𝑡 each phosphonic acid-2 propanol heterodimer ion. We then use van’t Hoff plots to determine ∆𝐻1,𝑆and ∆𝑆1,𝑆. Second, we apply Equations (10) and (11) to obtain ∆𝐻1,𝑟𝑒𝑓 and ∆𝑆1,𝑟𝑒𝑓. Ultimately this method yields two types of thermodynamic properties, one corresponding to the value at saturation (subscript “S”) and with respect to the reference temperature 298 K (subscript “ref”). For clarity, and to avoid a moving reference state with respect to temperature, we use the constant reference state conditions to discuss whether proton bound heterodimer formation is enthalpically or entropically favored (as is common convention). We also remark that Equation (7) can be used to generate separate van’t Hoff plots to infer ∆𝐻1,𝑟𝑒𝑓 and ∆𝑆1,𝑟𝑒𝑓; doing so yields equivalent results to the approach adopted here. We also note that Equation (1) can be applied only in the instance where a single modifier molecule incorporated to form a proton bound heterodimer. Though this is a clear restriction of the model, we believe this is a reasonable assumption (and examined its 𝑃

validity for the results obtained), as the ratio (𝑃 (𝑇)) remains comparatively small in this 𝑠𝑎𝑡 study. In Equation (1) all variables are experimentally determined directly, except the change in Gibbs free energy, and the collision cross section ratio, Ω1/Ω0. In our prior study,21 Ω1/Ω0 was estimated using molecular modeling predictions.40 However, careful consideration of the experimental conditions and protocol suggest that this ratio can instead be estimated from the experimental data provided the shift in drift time is brought about only by transient heterodimer formation; in this instance the maximum drift time shift would be achieved when the heterodimer remains bound together throughout the


Page 11 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


mobility measurement. Therefore, Ω1 may be safely estimated from the experimental data plot where the drift time ratio remains unchanged with any further increase in the 𝑃

saturation ratio (𝑃 (𝑇)). Interestingly, this experimental approach obviates the need for 𝑠𝑎𝑡 molecular modeling to obtain the ratio of Ω1/Ω0. To aid in estimating the impact of this experimentally determined fitting parameter, the collision cross section ratio Ω1/Ω0, was calculated with Ω1 inferred from measurements at 80, 120, and 160 µL/hr respectively, in order to fit the experimental results to Equation 1. With the collision cross section ratio defined, ∆𝐺1,𝑆 is the only remaining free parameter in Equation 1 and can be directly fit to experimental data.

RESULTS AND DISCUSSION Using a homologous set of chemical warfare agent degradation products simulants (phosphonic acids), the goal of the present study is to quantify the changes in enthalpy and entropy associated with the formation of proton bound heterodimer ions of the drift gas modifier (2-propanol) and the target analytes. Using this method focused on the relative shift in ion mobility (not simply changes in intensity) we demonstrate that variable temperature IM-MS measurements with vapor modifiers can yield the thermodynamic values for proton-bound heterodimer ions. In the absence of 2-propanol the positive ion mode reduced mobilities (i.e. the mobilities corrected to 273.15 K and 760 Torr, through linear scaling, as shown in the supplemental information SI Equation S6) recorded experimentally for MPA, EPA, and PPA were 1.98 cm2V-1s-1, 1.89 cm2V-1s-1 and 1.80 cm2V-1s-1 respectively, at 175 oC. Both the experimental and literature reduced mobilities of the target analyte ions are shown in Table 1. An apparent discrepancy exists between the only reported literature values for the target analyte ions and the values measured here.41 However, the temperature for the reported reduced mobility values in Table 1 was not stated in the publication and hence cannot be directly compared to our experimental measurements. It is known that the standard mobility reduction practice in



Page 12 of 32

IMS does not fully correct for temperature influences on mobility (i.e. it not does correct for the influence of the gas thermal speed, and how the effects of ion-induced dipole potentials affect measurements).1,42,43 At the same time, the reduced mobility values when examined from 100 oC to 175 oC for the target analyte ions varied almost linearly, from 1.85 +/0.01 to 1.98+/-0.01 cm2 V-1s-1, 1.79+/-0.01 to 1.89+/0.01 cm2 V-1s-1, and 1.70+/0.01 to 1.80+/0.01 cm2 V-1s-1 for MPA, EPA, and PPA ions respectively. The prior measurements reveal mobility values bounded by the lower and upper limits measured here however, at an apparent higher measurement temperature, and it appears that mobility reduction practices alone cannot explain the discrepancy between our measurements and prior work. However, as the main goal of the present study is to examine a relative shift in arrival time (as opposed to absolute arrival times), we elect not to examine this discrepancy further, and only note it for completeness. Figure 1, displaying plots of the relative total ion current as a function of drift time, demonstrates how vapor induced mobility shifts trend along the homologous series of phosphonic acids. The bottom traces exhibit peaks for each target analyte ion in pure nitrogen gas without the modifier. These peaks correspond to drift times of 7.08 +/-0.02 ms, 7.36+/-0.01 ms and 7.80+/-0.01 ms for MPA, EPA, and PPA respectively. With the introduction of the drift gas modifier (2-propanol) at 20 µL/hr, the drift time of MPA, EPA and PPA increased respectively to 8.71 +/- 0.01 ms, 8.83 +/- 0.01 ms, and 9.27 +/- 0.01 ms. When examining the mass spectra with the modifier added, at no time was a complex with more than one 2-propanol bound observed in the m/z domain and there was no appreciable Faraday plate ion signal at drift times that could reasonably correspond to proton bound heterotrimer or higher order cluster. Additionally, the drift time for each bare protonated phosphonic acid ion was identical to the drift time for the proton bound heterodimers from the mass-selected mobility spectra obtained from the frequency modulated IMS experiment (as seen in Figure S7). These observations strongly support the vapor modifier theory utilized subsequently in this work, which is based upon the assumptions that (1) mobility shifts at low relative vapor concentration


Page 13 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


are based upon the transient formation of a heterodimer complex during mobility measurement, and (2) measurement times are long enough for the bare ion to achieve equilibrium with the proton bound heterodimer. Because the heterodimer can be (and likely is) weakly bound, it may and likely does dissociate upon introduction to the mass spectrometer. However, such dissociation does not influence data analysis (as it would in the case of high pressure mass spectrometry), as the analysis is based solely upon the shift in drift time observed in the presence of the vapor modifier. When present, the modifier induced a comparatively greater relative and absolute shift in drift time for MPA compared to the rest of the target analytes. This is largely because the MPA ion is physically the smallest examined; while smaller ion size can enhance the influence of the positive charge and promote binding of the polar modifier molecule most directly,44-45 the ratio Ω1/Ω0 will be larger for smaller ions and is not dependent upon the thermodynamic parameters for heterodimer formation. We remark that the asymptotic relative shift in mobility is directly related to Ω1/Ω0. Of greater interest is inference of the Gibbs free energy, which can only be deduced from measurements at multiple drift modifier vapor concentrations. The mobility with the modifier (Kn) and without the modifier (Kn0) and the drift time ratio are plotted in Figure 2 against the saturation ratio( 𝑃 of the modifier (2-propanol) for MPA, EPA, and PPA at 175 0C. At the maximum 𝑃𝑠𝑎𝑡(𝑇))

modifier flow rate (160 µL/hr), the reduced mobilities recorded were 1.47 cm2 V-1s-1, 1.45 cm2 V-1s-1, 1.40 cm2 V-1s-1, corresponding to mobility decreases of 25.8 %, 23.2%, 22.2% for MPA, EPA, and PPA ions respectively. Though these shifts are appreciable, they are well within the expected extent for the binding of a single modifier molecule and the formation of a heterodimer only. Across the examined 2-propanol saturation ratio range, the mobility ratio increases across the homologous series of the phosphonic acid with MPA ions experiencing the lowest mobility ratio while PPA experienced the highest 𝑃

change in mobility ratio. For the lowest saturation ratio (𝑃 (𝑇)), there is a sharp change 𝑠𝑎𝑡



Page 14 of 32

𝑃

in drift time ratios and mobility ratios while at higher saturation ratios (𝑃 (𝑇)), the relative 𝑠𝑎𝑡 drift time ratio and the mobility ratios remains somewhat constant. This too agrees with model expectations, in that if only a single vapor molecule binds to the analyte ion forming a heterodimer transiently, the drift time/mobility ratios will reach an asymptotic 𝛺1𝜇1/2 0 value of ( ). Estimations of this ratio from the mobility ratios at the highest vapor 𝛺0𝜇1/2 1

modifier concentrations at 175 0C were 1.34, 1.33, 1.30 for MPA, 1.31, 1.29, 1.27 for EPA and 1.29, 1.27, 1.25 for PPA using vapor modifier introduction rates of 160, 120 and 80 µL/hr respectively. The estimated 2-propanol-protonated phosphonic acid collision cross sections (Ω1) (for 160 µL/hr liquid feed flow rate of the modifier) are reported in Table 2. To obtain the ∆𝐺1,𝑠, the Gibbs free energy of association at saturation, at any given temperature, the mean square error between the relative drift time ratios and the predictions from Equation 1 was minimized. Equation (1) based best fit curves from this method are plotted in Figure 2. ∆𝐺1,𝑠 and ∆𝐺1,𝑟𝑒𝑓 (Equation 7) inferred from the plotted curves at 80, 120, and 160 µL/hr (as seen in SI Table S3 and S4 respectively) were averaged and summarized in Tables 3 and 4, respectively. Additionally, Table S1 lists the natural logarithm of the equilibrium constants (ln(K)) determined from ∆𝐺1,𝑠 and Table S2 lists the natural logarithm of the equilibrium constants determined from ∆𝐺1,𝑟𝑒𝑓. While equilibrium constant logarithms increase with increasing temperature when defined based upon ∆𝐺1,𝑠, they decrease with increasing temperature based upon ∆𝐺1,𝑟𝑒𝑓, which suggests that heterodimer formation is exothermic. The absolute values of the reported average Gibbs free energy in the Tables 3 and 4 vary slightly for each target analyte ion at any given measurement temperature; however, they are not statistically different from one another. The similarity between ions suggests that heterodimer formation via the binding of 2-propanol to the target analyte ions may be largely controlled by the chemical functional group participating in heterodimer formation for the examined ions.


Page 15 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


After the inferred ∆𝐺1,𝑠 values were extracted at each temperature, van’t Hoff plots were readily obtained, and are shown in Figure 3. Such plots yield the enthalpies and entropies for heterodimer formation from the ion and modifier. After conversion to a standard reference state using Equations (10) and (11), we find the that ∆𝐻1,𝑟𝑒𝑓 values for heterodimer ions formed from MPA, EPA and PPA were -2.05 kcal/mol, -0.987 kcal/mol, -1.68 kcal/mol, respectively; the ∆𝑆1,𝑟𝑒𝑓 values for the same heterodimers were 13.7 cal/molK, 15.9 cal/molK, and 14.01 cal/molK, respectively. It is important to note that we base our analysis on the binding of a single modifier molecule. However, at higher 𝑃

saturation ratio (𝑃 (𝑇)), the experimental curve begins to deviate relative to the 𝑠𝑎𝑡 clustering model curve which occurs after approximately 80 µL/hr of the liquid modifier is introduced into the system (See Figure 2). Close comparisons between the experimental and model results illustrate an upward trend in the experimental curve relative to the clustering model. Though not shown, extension of the model to include a second vapor molecule binding event was considered, as the binding of additional vapor molecules would lead to additional mobility shifts and is likely the source of this deviation. When applying the appropriate fit constraints (i.e. Ω1 0.92). Data for the plots shown here were collected at 175 oC and 150 oC.


Page 21 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


FIGURE 3. van’t Hoff plot comparing the natural log of the equilibrium constant as a function of inverse of temperature for the clustering of 2-propanol with the target analyte ion. The inferred equilibrium constants (calculated from Gibbs free energy (∆𝐺1,𝑆) inferred at variable saturation state) for 80, 120, and 160 µL/hr were averaged and plotted with error bars.



Page 22 of 32

TABLE 1 Summary of molecular structure, molecular weights and reduced mobilities for the target analyte ions TARGET

MOLECULAR

ANALYTE

WEIGHT

STRUCTURE

ION

Ko (cm2/Vs) in N2

(g/mol)

Ko (cm2/Vs) in N2 Drift Gas

Drift Gas

EXPERIMENTALb

LITERATUREa Methyl

96.022

[MPA+H]+

1.88

1.98 +/- 0.02

110.049

[EPA+H]+

1.81

1.89 +/- 0.02

124.076

[PPA+H]+

1.72

1.80 +/- 0.02

phosphonic acid (MPA)

Ethyl phosphonic acid (EPA)

Propyl phosphonic acid (PPA) a

Temperature of the experiment not stated

b

Temperature of the experiment was done at 175 0C. Replicates measurements were

made as outlined in the experimental section a standard deviation of 0 in the hundredths place. However, the errors reported are sources of error related to the units used to control instrument variables (i.e. power supplies, drift length measurements and temperature controllers)


Page 23 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60




Page 24 of 32

TABLE 2. Summary of the Measured collision cross section (CCS) as well as Reduced Masses (in Nitrogen) of the ions of interest at 175 0C. The CCS for the ion-neutral vapor molecule complex were reported at 160 uL/hr liquid modifier flow rate. ION

experimental CCS/ A2

reduced mass (Da)

Methylphosphonic acid-H+

95 +/- 2

21.73

Methylphosphonic acid-2-propanol-H+

127 +/- 2

23.76

Ethylphosphonic acid-H+

98 +/- 2

22.36

Ethylphosphonic acid-2-propanol-H+

128 +/- 2

24.06

Propylphosphonic acid-H+

102 +/- 2

22.88

Propylphosphonic acid-2-propanol-H+

131 +/- 2

24.32


Page 25 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


TABLE 3. Summary of Gibbs free energies at variable saturation state (∆𝑮𝟏,𝑺) at any given measurement temperature Temperature ∆𝐺1,𝑆 (kcal/mol) (K) MPA

∆𝐺1,𝑆

∆𝐺1,𝑆

(kcal/mol)

(kcal/mol)

EPA

PPA

448.15

-13.62

-13.70

-13.49

423.15

-12.05

-12.13

398.15

-11.05

373.15

-9.71

Compound

∆𝐻1,𝑆

∆𝑆1,𝑆

(kcal/mol)

(cal/K.mol)

MPA

8.81

50.10

-12.10

EPA

9.87

52.4

-10.98

-10.87

PPA

9.18

50.5

-9.70

-9.68

TABLE 3. Gibbs free energy (∆𝐺1,𝑆) at variable saturation state at any given measurement temperature. The inferred Gibbs free energy values and log of equilibrium constants from 80, 120, and 160 µL/hr flow rate (liquid modifier feed) were averaged. ∆𝐻1,𝑆 and ∆𝑆1,𝑆 were calculated from the slope and intercept respectively in Figure 3.



Page 26 of 32

TABLE 4. Summary of Gibbs free energies inferred at a consistent reference state (∆𝑮𝟏,𝒓𝒆𝒇) at any given measurement temperature Temperature ∆𝐺1,𝑟𝑒𝑓 (kcal/mol) (K) MPA

∆𝐺1,𝑟𝑒𝑓 (kcal/mol)

448.15

-8.24

-8.25

-8.03

MPA

-2.05

13.7

423.15

-7.71

-7.58

-7.55

EPA

-0.987

15.9

398.15

-7.56

-7.35

-7.23

PPA

-1.68

14.1

373.15

-7.14

-6.97

-6.96

EPA

∆𝐺1,𝑟𝑒𝑓

Compound

(kcal/mol)

∆𝐻1,𝑟𝑒𝑓 (kcal/mol)

∆𝑆1,𝑟𝑒𝑓 (cal/K mol)

PPA

TABLE 4. Gibbs free energy (∆𝐺1,𝑟𝑒𝑓) inferred from the experiment through Equation 7 at any given measurement temperature. The inferred Gibbs free energy values and log of equilibrium constants from 80, 120, and 160 µL/hr flow rate (liquid modifier feed) were averaged. 𝑙𝑛𝐾1,𝑟𝑒𝑓 calculated from ∆𝐺1,𝑟𝑒𝑓 was used to create a van’t Hoff to yield ∆𝐻1,𝑟𝑒𝑓 and ∆𝑆1,𝑟𝑒𝑓 from the slope and intercept respectively.


Page 27 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


References (1) Siems, W. F.; Viehland, L. A.; Hill, H. H., Jr. Improved Momentum-Transfer Theory for Ion Mobility. 1. Derivation of the Fundamental Equation. Anal. Chem. 2012, 84 (22), 9782–9791. (2) Albritton, D. L.; Miller, T. M.; Martin, D. W.; McDaniel, E. W. Mobilities of MassIdentified H 3+ and H+ Ions in Hydrogen. Physical Review 1968, 171 (1), 94. (3) Revercomb, H. E.; Mason, E. A. Theory of Plasma Chromatography/gaseous Electrophoresis. Review. Anal. Chem. 1975, 47 (7), 970–983. (4) Gabelica, V.; Shvartsburg, A. A.; Afonso, C.; Barran, P. E.; Benesch, J. L. P.; Bleiholder, C.; Bowers, M. T.; Bilbao, A.; Bush, M. F.; Campbell, J. L.; et al. Recommendations for Reporting Ion Mobility Mass Spectrometry Measurements. 2018. (5) Reinecke, T.; Clowers, B. H. Implementation of a Flexible, Open-Source Platform for Ion Mobility Spectrometry. HardwareX 2018, 4, e00030. (6) Asbury, G. R.; Klasmeier, J.; Hill, H. H., Jr. Analysis of Explosives Using Electrospray Ionization/ion Mobility Spectrometry (ESI/IMS). Talanta 2000, 50 (6), 1291–1298. (7) Steiner, W. E.; Klopsch, S. J.; English, W. A.; Clowers, B. H.; Hill, H. H. Detection of a Chemical Warfare Agent Simulant in Various Aerosol Matrixes by Ion Mobility Time-of-Flight Mass Spectrometry. Anal. Chem. 2005, 77 (15), 4792–4799. (8) Mäkinen, M.; Nousiainen, M.; Sillanpää, M. Ion Spectrometric Detection Technologies for Ultra-Traces of Explosives: A Review. Mass Spectrom. Rev. 2011, 30 (5), 940–973. (9) Karasek, F. W.; Hill, H. H., Jr; Kim, S. H. Plasma Chromatography of Heroin and Cocaine with Mass-Identified Mobility Spectra. J. Chromatogr. 1976, 117 (2), 327– 336. (10) Shi, L.; Holliday, A. E.; Shi, H.; Zhu, F.; Ewing, M. A.; Russell, D. H.; Clemmer, D. E. Characterizing Intermediates along the Transition from Polyproline I to Polyproline II Using Ion Mobility Spectrometry-Mass Spectrometry. J. Am. Chem. Soc. 2014, 136 (36), 12702–12711. (11) Pringle, S. D.; Giles, K.; Wildgoose, J. L.; Williams, J. P.; Slade, S. E.; Thalassinos, K.; Bateman, R. H.; Bowers, M. T.; Scrivens, J. H. An Investigation of the Mobility Separation of Some Peptide and Protein Ions Using a New Hybrid Quadrupole/travelling Wave IMS/oa-ToF Instrument. Int. J. Mass Spectrom. 2007, 261 (1), 1–12. (12) Larriba, C.; Hogan, C. J., Jr. Ion Mobilities in Diatomic Gases: Measurement versus Prediction with Non-Specular Scattering Models. J. Phys. Chem. A 2013, 117 (19), 3887–3901. (13) Shibue, M.; Mant, C. T.; Hodges, R. S. Effect of Anionic Ion-Pairing Reagent Hydrophobicity on Selectivity of Peptide Separations by Reversed-Phase Liquid



Page 28 of 32

Chromatography. J. Chromatogr. A 2005, 1080 (1), 68–75. (14) Shibue, M.; Mant, C. T.; Hodges, R. S. Effect of Anionic Ion-Pairing Reagent Concentration (1–60mM) on Reversed-Phase Liquid Chromatography Elution Behaviour of Peptides. J. Chromatogr. A 2005, 1080 (1), 58–67. (15) Li, C.; Hogan; Jr.; J., C. Vapor Specific Extents of Uptake by Nanometer Scale Charged Particles. Aerosol Sci. Technol. 2017, 51 (5), 653–664. (16) Fernandez-Maestre, R. Buffer Gas Additives (modifiers/shift Reagents) in Ion Mobility Spectrometry: Applications, Predictions of Mobility Shifts, and Influence of Interaction Energy and Structure. J. Mass Spectrom. 2018, 53 (7), 598–613. (17) Roscioli, K. M.; Lamabadusuriya, M. R.; Harden, C. S.; Midey, A. J.; Wu, C.; Siems, W. F.; Hill, H. H. Structure Selective Ion Molecule Interactions (SSIMI) in Ion Mobility Spectrometry. Int. J. Ion Mobil. Spectrom. 2014, 17 (1), 43–53. (18) Fernández-Maestre, R.; Wu, C.; Others. Ammonia as a Modifier in Ion Mobility Spectrometry: Effects on Ion Mobilities and Potential as a Separation Tool. J. Chil. Chem. Soc. 2014, 59 (1), 2398–2403. (19) Rawat, V. K.; Vidal-de-Miguel, G.; Hogan, C. J., Jr. Modeling Vapor Uptake Induced Mobility Shifts in Peptide Ions Observed with Transversal Modulation Ion Mobility Spectrometry-Mass Spectrometry. Analyst 2015, 140 (20), 6945–6954. (20) Oberreit, D.; Rawat, V. K.; Larriba-Andaluz, C.; Ouyang, H.; McMurry, P. H.; Hogan, C. J., Jr. Analysis of Heterogeneous Water Vapor Uptake by Metal Iodide Cluster Ions via Differential Mobility Analysis-Mass Spectrometry. J. Chem. Phys. 2015, 143 (10), 104204. (21) Kwantwi-Barima, P.; Ouyang, H.; Hogan, C. J., Jr; Clowers, B. H. Tuning Mobility Separation Factors of Chemical Warfare Agent Degradation Products via Selective Ion-Neutral Clustering. Anal. Chem. 2017, 89 (22), 12416–12424. (22) Kebarle, P.; Hogg, A. M. Heats of Hydration and Solvation by Mass Spectrometry. J. Chem. Phys. 1965, 42 (2), 798–799. (23) Staley, R. H.; Beauchamp, J. L. Intrinsic Acid-Base Properties of Molecules. Binding Energies of lithium(1+) Ion to .pi.- and N-Donor Bases. J. Am. Chem. Soc. 1975, 97 (20), 5920–5921. (24) Ibrahim, Y.; Alsharaeh, E.; Dias, K.; Meot-Ner, M.; El-Shall, M. S. Stepwise Hydration and Multibody Deprotonation with Steep Negative Temperature Dependence in the Benzene•+- Water System. J. Am. Chem. Soc. 2004, 126 (40), 12766–12767. (25) Ibrahim, Y.; Alsharaeh, E.; Meot-Ner, M. MS El-Shall, Scheiner, SJ Am. In Chem. Soc; 2005; Vol. 127, p 7053. (26) Tang, I. N.; Lian, M. S.; Castleman, A. W. Mass Spectrometric Study of Gas‐phase Clustering Reactions: Hydration of the Monovalent Strontium Ion. J. Chem. Phys. 1976, 65 (10), 4022–4027. (27) Hoyau, S.; Norrman, K.; McMahon, T. B.; Ohanessian, G. A Quantitative Basis for a Scale of Na+ Affinities of Organic and Small Biological Molecules in the Gas Phase. J. Am. Chem. Soc. 1999, 121 (38), 8864–8875. (28) Castleman, A. W.; Holland, P. M.; Lindsay, D. M.; Peterson, K. I. The Properties


Page 29 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


of Clusters in the Gas Phase. 2. Ammonia about Metal Ions. J. Am. Chem. Soc. 1978, 100 (19), 6039–6045. (29) Castleman, A. W.; Peterson, K. I.; Upschulte, B. L.; Schelling, F. J. Energetics and Structure of Na+ Cluster Ions. Int. J. Mass Spectrom. Ion Phys. 1983, 47, 203–206. (30) Peterson, K. I.; Mark, T. D.; Keesee, R. G.; Castleman, A. W. Thermochemical Properties of Gas-Phase Mixed Clusters: Water/carbon Dioxide with sodium(1+). J. Phys. Chem. 1984, 88 (13), 2880–2883. (31) Guo, B. C.; Conklin, B. J.; Castleman, A. W. Thermochemical Properties of Ion Complexes Na+(M)n in the Gas Phase. J. Am. Chem. Soc. 1989, 111 (17), 6506– 6510. (32) Hamid, A. M.; Sharma, P.; El-Shall, M. S.; Hilal, R.; Elroby, S.; Aziz, S. G.; Alyoubi, A. O. Hydration of the Pyrimidine Radical Cation and Stepwise Solvation of Protonated Pyrimidine with Water, Methanol, and Acetonitrile. J. Chem. Phys. 2013, 139 (8), 084304. (33) Keesee, R. G.; Castleman, A. W. Thermochemical Data on Gas‐Phase Ion‐Molecule Association and Clustering Reactions. J. Phys. Chem. Ref. Data 1986, 15 (3), 1011–1071. (34) Meot-Ner, M.; Sieck, L. W. The Ionic Hydrogen Bond. 1. Sterically Hindered Bonds. Solvation and Clustering of Protonated Amines and Pyridines. J. Am. Chem. Soc. 1983, 105 (10), 2956–2961. (35) Meot-Ner, M.; Field, F. H. Association and Solvation Reactions of Protonated Gaseous Amino Acids. J. Am. Chem. Soc. 1974, 96 (10), 3168–3171. (36) Holland, P. M.; Castleman, A. W. The Thermochemical Properties of Gas‐phase Transition Metal Ion Complexes. J. Chem. Phys. 1982, 76 (8), 4195–4205. (37) Adams, N. G.; Smith, D. The Selected Ion Flow Tube (SIFT); A Technique for Studying Ion-Neutral Reactions. Int. J. Mass Spectrom. Ion Phys. 1976, 21 (3), 349–359. (38) Smith, D.; Adams, N. G. The Selected Ion Flow Tube (Sift): Studies of IonNeutral Reactions. In Advances in Atomic and Molecular Physics; Bates, D., Bederson, B., Eds.; Academic Press, 1988; Vol. 24, pp 1–49. (39) Morrison, K. A.; Siems, W. F.; Clowers, B. H. Augmenting Ion Trap Mass Spectrometers Using a Frequency Modulated Drift Tube Ion Mobility Spectrometer. Anal. Chem. 2016, 88 (6), 3121–3129. (40) Shrivastav, V.; Nahin, M.; Hogan, C. J.; Larriba-Andaluz, C. Benchmark Comparison for a Multi-Processing Ion Mobility Calculator in the Free Molecular Regime. J. Am. Soc. Mass Spectrom. 2017, 28 (8), 1540–1551. (41) Asbury, G. R.; Wu, C.; Siems, W. F.; Hill, H. H. Separation and Identification of Some Chemical Warfare Degradation Products Using Electrospray High Resolution Ion Mobility Spectrometry with Mass Selected Detection. Anal. Chim. Acta 2000, 404 (2), 273–283. (42) Viehland, L. A.; Siems, W. F. Uniform Moment Theory for Charged Particle Motion in Gases. J. Am. Soc. Mass Spectrom. 2012, 23 (11), 1841–1854.



Page 30 of 32

(43)

Tabrizchi, M.; Rouholahnejad, F. Comparing the Effect of Pressure and Temperature on Ion Mobilities. J. Phys. D Appl. Phys. 2005, 38 (6), 857. (44) Nasibulin, A. G.; de la Mora, J. F.; Kauppinen, E. I. Ion-Induced Nucleation of Dibutyl Phthalate Vapors on Spherical and Nonspherical Singly and Multiply Charged Polyethylene Glycol Ions. J. Phys. Chem. A 2008, 112 (6), 1133–1138. (45) Tauber, C.; Chen, X.; Wagner, P. E.; Winkler, P. M.; Hogan, C. J.; Maisser, A. Heterogeneous Nucleation onto Monoatomic Ions: Support for the KelvinThomson Theory. Chemphyschem 2018. (46) Köhler, H. The Nucleus in and the Growth of Hygroscopic Droplets. Transactions of the Faraday Society 1936, 32, 1152–1161. (47) Gamero-Castano, M.; de la Mora, J. F. Ion-Induced Nucleation: Measurement of the Effect of Embryo’s Size and Charge State on the Critical Supersaturation. J. Chem. Phys. 2002, 117 (7), 3345–3353.


Page 31 of 32 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


TOC GRAPHIC



Accounting of vapor-induced mobility shifts yields thermodynamic values for weakly-bound, proton-bound heterodimers.

ACS Paragon Plus Environment

Page 32 of 32

Deducing Proton-Bound Heterodimer Association Energies from Shifts

Recommend Documents