From ions in solution to ions in the gas phase - the mechanism of

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Paul Kebarle and Liang Tang Department of Chemistry University of Alberta Edmonton, Alberta Canada T6G 2G2

It is no exaggeration to say that close to half of inorganic chemistry, organic chemistry, and biochemistry involves ions in solution. A technique t h a t allows ions to be transferred from solution to the gas phase and subjected to m a s s spectrometric analysis should therefore be of enormous importance. Electrospray (ES) is such a technique. The transfer of ions from the gas phase to solution is a natural process. In the presence of solvent molecules such as H,O, naked gas-phase ions such as H30+ or Na+ will spontaneously form ion-solvent molecule clusters such a s H,O+(H,O), a n d Na+(H,O),. If the pressure of the solvent vapor is somewhat above the s a t u r a t i o n vapor pressure, these clusters will grow to small droplets. To visualize this process one needs only to t h i n k of t h e Wilson cloud chamber, where such a process occurs in front of one’s eyes. In the past 30 years mass spectro972 A

metric studies of ion-molecule clusters in the gas phase (I)have contributed much to the understanding of ion-solvent molecule interactions in solution. It is much easier to und e r s t a n d t h e n a t u r e of t h e ionsolvent bonding when the ion interacts with only one or a few solvent molecules, and it is these initial interactions t h a t dominate ion solvation in solution (1). The transfer of ions from solution to the gas phase is a desolvation process, and thus it is strongly “unnatural” or endoergic. The free energy required when a mole of Na’ is transferred from aqueous solution to the gas phase is very large (2) Na’ (aq) + Na’ (g) -AGLl(Na+) = 98.2 kcaYmol Analytical m a s s spectrometric methods in which ions a r e “transferred” from solution to t h e g a s phase, such a s fast atom bombardment (FAB), plasma desorption, and laser desorption, existed before the introduction of electrospray mass spectrometry (ESMS). The energy required for ion transfer to the gas phase in these earlier methods is supplied by complex high - energy collision cascades and highly localized heating, resulting in additional processes such as net ionization (creation of ions from neutrals) and fragmentation of ions. ESMS, in contrast, is the

ANALYTICAL CHEMISTRY, VOL. 65, NO. 22, NOVEMBER 15, 1993

“purest” form of transfer of ions from solution to the gas phase; little if any extra internal energy is imparted to the ions. ES affords ion transfer from solution to t h e gas phase for a large range of ion types: the simple singly charged electrolytes such as Na’ and C1- and organic protonated bases BH+ and X-;group I1 ions such as Ca”, SP’, and Ba2+,as well as doubly a n d triply charged transition metal and lanthanide ions and complexes thereof; bioorganic ions such as multiply- protonated peptides and proteins of molecular mass 100,000 Da; and multiply-deprotonated, negatively charged nucleic acids. ESMS was introduced by Yamashita and Fenn in 1984 (31, and it took a few years before the importance of the method was recognized (4-6). Rec e n t advances (7-9)confirm t h e widely held view t h a t ES has ushered in a new era of the mass spectrometric analysis of biomolecules. The exciting applications of ESMS have also created great interest in t h e mechanism by which t h e gasphase ions required for t h e mass spectrometric analysis are produced by the ES device. Ideally, a n understanding of the mechanism should not only provide aesthetic satisfaction, it should be of significant practical use. Understanding the mechanism essentially provides a mental map that allows planning of experi0003- 2700/93/0365-972A/$04.00/0 0 1993 American Chemical Society

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The ES mechanism: Features of importance to ESMS There a r e four major processes in ESMS: t h e production of charged droplets from electrolyte dissolved in a solvent; shrinkage of charged droplets by solvent evaporation and repeated droplet disintegrations (fis sions), leading ultimately to very small, highly charged droplets capable of producing gas-phase ions; the mechanism of gas -phase ion produc tion; and secondary processes, by which gas - phase ions are modified in the atmospheric and the ion sampling regions of the spectrometer. Although this last topic is important, it includes a wide area of phenomena,

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0 position to use the full range of their background and imagination to develop new applications of the ESMS technique. ES existed long before its application to MS. I t is a method of considerable importance for t h e electros t a t i c dispersion of liquids. The interesting history and notable re search advances are very well described in Bailey's book, Electrostatic Spraying of Liquids (9).This research also provides the basis for the ESMS mechanistic studies. The new aspect in the mass spectrometric application concerns the production of gas phase ions from the ions in the solution; this was of no interest to the spray researchers (9).

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ments and making rational choices of experimental parameters leading to a desired result. Unfortunately, the understanding of the ES mechanism is not yet complete. Nevertheless, t h e whys a n d why nots for many experimental parameters are understood. What is still uncertain is exactly how the gas-phase ions are produced from the very small and highly charged droplets. Readers of this REPORT will learn about current ideas but will not find certainty. However, much useful and interest ing information is obtained about the conditions existing in ESMS when one attempts to pin down the mechanism involved. Fortunately, for practical purposes, the predictions of the two presently favored mechanisms turn out to be quite similar. ESMS has a great advantage over the other methods because the ionic solutions used for transfer of ions to the gas phase are the same solutions used in conventional wet chemistry. No unusual liquid matrices such as the glycerol used in FAB and matrixassisted laser desorption ionization are required. Now that good ESMS interfaces are available on commer cia1 mass spectrometers, the users and not the interface developers are the important players. These users a r e typically workers i n solution chemistry. Once these individuals have acquired some understanding of the mechanism, they will be in a

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some of which are not well understood. Therefore these effects will not be considered here. Production of charged droplets at the ES capillary tip As shown in the schematic representation of the ES events in air at atmospheric pressure (Figure l), a voltage V, of 2-3 kV is applied to the metal capillary, which is typically 0.2 mm 0.d. and 0.1 mm i.d. and located 1-3 cm from the large planar counter electrode. I n ESMS this counter electrode has an orifice leading to the mass spectrometric sampling system. Because the capillary tip is very narrow, the electric field E in the air at the capillary tip is very high (E = lo6 V/m). When the capillary of radius r, is located at a distance d from the planar counter electrode, the magnitude of E, for a given potential V , is given by (10,11) Ec = 2 VChc1n(4dlrc) Equation 2 provides the field at the capillary tip in the absence of solution. The field E, is proportional to V,, and the most important geometry parameter is r,. E, is essentially inversely proportional t o r,; it d e creases slowly, with a logarithmic dependence, with the distance d. For example, when V, = 2000 V, rc = 0.2 mm, a n d d = 2 cm, E, = 3.3 x lo6 V/m.


9

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REPOR7 The typical solution used in ESMS consists of a dipolar solvent in which electrolytes a r e at least somewhat soluble. We use methanol as the solvent and NaCl or BHCl (where B is an organic nitrogen base) as the electrolytes. Essentially, these electrolytes are totally dissociated into Na' and C1- as well as BH' and C1- in the solvent. Low electrolyte concentrations, routinely in the M range, are required for the operation of E$. The imposed field E will also partially penetrate the liquid at the capillary tip. To simplify the discussion, we will assume that all experiments are carried out in the positive ion mode, which has a positive capillary potential. When the capillary is the positive electrode, some positive ions in the liquid will drift toward the liquid surface and some negative ions drift away from it until the imposed field inside the liquid is essentially removed by this charge redistribution. However, the accumulated positive charge at the surface leads t o destabilization of the surface because the positive ions are drawn downfield but cannot escape from the liquid. The surface is drawn out downfield such t h a t a liquid cone forms. This is called a Taylor cone, after the researcher who was one of the first to investigate the conditions under which a stable liquid cone can exist with the competing forces of an electric field and the surface tension of the liquid (12). At a sufficiently high field E, the cone is not stable and a liquid filament with a diameter of a few micrometers, whose surface is enriched on positive ions, is emitted from the Taylor cone tip. At some distance downstream, the liquid filament becomes unstable and forms separate droplets. The droplets' surfaces are enriched with positive ions for which there are no negative counterions in t h e droplet (i.e., t h e droplets a r e charged with a n excess of positive electrolyte ions). The length of the unbroken liquid filament decreases if t h e field E is increased. At higher fields, a multispray condition is reached in which the central cone disappears and droplet emission occurs from a crown of four to six short liquid tips formed a t the rim of the capillary (13). We came to the conclusion that the ion separation described above is the mechanism responsible for droplet charging, largely on t h e basis of mass spectra observed with electrolytes in the 10-5-10-3 M range (13, 14). The positive and negative ions 974 A

Figure 1. Schematic representation of processes in ESMS. The very high electric field imposed by the power supply causes an enrichment of positive electrolyte ions at the meniscus of the solution at the metal capillary tip. This net charge is pulled downfield, expanding the meniscus into a cone that emits a fine mist of positively charged droplets. Solvent evaporation reduces the volume of the droplets at constant charge, leading to fission of the droplets. Charge balance is attained in the ES device by electrochemical oxidation at the positive electrode and reduction at the negative electrode.

observed in the spectra were always the positive and negative ions of the electrolytes present in the solution. Extraneous ions were observed only at high capillary voltages where electric (corona) discharges were occurring a t the capillary (15, 16). The ion separation mechanism, which is called the electrophoretic mechanism, is also most plausible on energetic grounds. Electrical double layers a r e already formed at low fields in electrolyte solutions. The resulting positive-negative ion redistribution reduces or completely removes the imposed field and therefore suppresses other forms of ionization such as ionization by electron removal from molecules (field ioniza tion), which requires very high electric fields. Another way to show t h a t ES results from electrophoretic charging is to deionize the solvent. Experiments involving methanol deionized by dis tillation, which reduced the conductivity from 10-6Q-1cm-1 (reagent grade methanol, M in i m p u r i t y electrolyte) to io-'^-^ cm-', showed t h a t the ES current decreased and became intermittent. The intensity of the observed impurity ions (NH',, Na') present in reagent -grade methanol became much

ANALYTICAL CHEMISTRY, VOL. 65, NO. 22, NOVEMBER 15,1993

lower in the purified solvent, and the signal fluctuated widely on a time scale of seconds (see Figure 3 in Reference 14). If the charge separation is electrophoretic, at a steady state the positive droplet emission will continuo u s l y c a r r y off p o s i t i v e i o n s . Considering t h e requirements for charge balance in such a continuous electric current device, and that only electrons can flow through the metal wire supplying the electric potential t o t h e electrodes (Figure l ) , one comes to the conclusion that the ES process should involve a n electro chemical conversion of ions to electrons. In other words, the ES device can be viewed a s a special type of electrolytic cell in which that part of t h e ion t r a n s p o r t does not occur through uninterrupted solution, but as charged droplets and later as ions in the gas phase. Therefore, a conventional electrochemical oxidation reaction should be occurring at the liquid-metal interface of the capillary. This reaction should be supplying positive ions to the solution by converting metal atoms to positive ions and electrons or by converting negative ions from the solution to neutral molecules and a n electron. This conversion process is accom-

plished by oxidation reactions such as M (s)+ M2'(aq) + 2e4 OH-(aq) + O,(g) + 2H20 + 4e- (3) assuming that an aqueous solution is electrosprayed. The actual lowest oxidation potential reaction expected to occur will depend entirely on the solvent and composition of the solution used. Proof for the occurrence of electrochemical oxidation a t the metal capillary was provided by Blades et al. (17). When a Zn capillary tip was used, the release of Zn2+to the solution could be detected with ESMS. Furthermore, the amount of Zn2+released to the solution corresponded to the amount required to carry the ES current. Similar results were observed with stainless steel capillaries that release Fe2+to the solution. The ions added to the solution by the electrolytic process are at a very low concentration, - 2 x lo-, M (17). They can be detected by MS, b u t they do not, in general, interfere with the detection of other ionic analytes present in the solution. The work reported by Van Berkel et al. (18)shows some very interesting results when dry nonprotic solvents are used. There is little explicit discussion in the pre - MS ES literature concerning the nature of the charge carriers in the droplets. Although this is a vital question to the mass spectrometrist, it is of limited interest in other applications of ES. Pfeifer and Hendricks (11)were probably the first authors who explicitly proposed and discussed the electrophoretic mechanism. More recently, Hayati, Bailey, and Tadros (19)also explicitly endorsed the electrophoretic charging mechanism in a comprehensive examination of features of the ES process. However, n e i t h e r research group considered the electrochemical nature of the process at the metalliquid interface. In a treatment of the cone instability and charged-droplet emission, Smith (20), who also assumed the electrophoretic mechanism of charging, provided a very useful equation for the electric field required for the

onset of charged-droplet emission (i.e., t h e onset of ES). When h i s equation is combined with Equation 2, which relates the field E, to the potential V, one obtains

VOn= 2 x 105(yrc)y21n(4d/rc) (4) where y is the surface tension of the solvent, r, is the capillary radius, and d is the distance between capillary tip and counter electrode (see Equation 2). The onset voltages V,, in Table I were evaluated with Equation 4 and illustrate the range of V,, values required by different solvents for d = 4 cm and r, = 0.1 mm. The surface of the solvent with the highest surface t e n s i o n ( H 2 0 ) i s t h e h a r d e s t to stretch out into a Taylor cone and a liquid filament; this leads to the requirement for the highest Von. Experimental verification of Equa tion 4 has been provided by Smith (20)and by work from our laboratory (15,16). For stable ES operation one needs to go a few h u n d r e d volts higher than Van. Using water as the solvent can lead to the initiation of an electric discharge from the capillary tip, particularly when the capillary is negative (i.e., in the negative ion mode 115,161).The occurrence of discharge is easily recognized in ESMS because it leads t o t h e a p pearance of discharge -produced ions. Thus, in the positive ion mode, the appearance of protonated solvent clusters such a s H30+(H20), from water and CH,OH; (CH,OH), from methanol is typical (15).These ions are abundantly produced by ES in the absence of discharge only when the solvent is acidified. The presence of discharge degrades t h e performance of ESMS. ES with pure water solvent is still possible when traces of discharge - suppressing gases such as SF, or larger quantities of 0, are added to the air surrounding the ES capillary (6, 15-1 7). The high potentials required (Table I) also illustrate that air a t atmospheric pressure is not only a conven i e n t b u t also a very s u i t a b l e ambient g a s for ES, particularly when solvents such as water are to be used. The 0, in air captures elec-

trons (although less efficiently than SF,) and has a high electrical breakdown potential relative to gases such as N,, Ar, Ne, and He. ES in air or other gases below atmospheric pressure will also be adversely affected by electrical discharge because the electrical breakdown potential decreases (i.e., a discharge starts a t a lower voltage when the pressure is decreased). The capillary current I (Figure l ) , which results from charged droplets leaving the capillary, is of interest because i t provides a quantitative measure of the excess positive electrolyte ions produced by the spray. This current is easily measured and is generally i n t h e 1 x A10 x A range. Higher currents generally result from an electric discharge (15,16). A theoretical derivation giving the dependence of I on experimental parameters has been proposed by Hendricks (11).However, the derivation is based on some unproven assumptions. The equations obtained are

I =AHV;E'o"

(54

I = AHV;E'(hLc)"

(5b)

where A, is a constant that can be evaluated (11)and depends on the dielectric constant and surface tension of the solvent. Equation 5 brings out the parameters whose functional dependence has been subjected to experimental tests (11,13).These are the flow rate (volume/time) V,, the electric field E a t the cone tip, and t h e conductivity of the solution 0 . Equation 5b expresses the conductivity CT in terms of the relationship 0

=a:c

(6)

which holds a t the low concentrations C of electrolytes prevailing in ESMS. The equivalent mglar conductivity of the electrolyte h , gives the dependence of t h e conductivity on the specific nature of the electrolyte ions. The exponents v = 0.57; E = 0.43, and n = 0.43associated with the flow rate, electric field, and the conductivity are predicted by the Hendricks e q u a t i o n ( 1I). E x p e r i m e n t a l l y the following values are found: v = 0.5, E = 0.5, and n = 0.2-0.3 (11, 14, 21). Because of the small exponents, the changes of I with flow rate, electric field, and concentration of the electrolyte are small. Experimentally determined changes of I with the flow rate V, are shown in Figure 2a. Willoughby e t al. have recently presented (22)a quantitative theory based on an analysis of the conver-


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REPOR7 sion of energy from electrostatic potential energy to surface and kinetic energy of the droplets that occurs at t h e liquid cone t i p a n d droplet stream. Although details were not available, the approach appears very promising and may provide a reliable and useful relationship between the import ant parameters involved. Comparison of the current I with the mass spectrometrically detected intensity I(BH+)of the analyte (Figures 2a and 2b) shows that the current I and the gas-phase ion current I(BH+) are not closely related. The absence of a close correlation between these two quantities underscores t h e need t o understand the process by which gas-phase ions are produced from the charged droplets. A more detailed general examination of this process will be given later.

Here we consider briefly the effect of flow rate Vf and the concentration C. As shown below, the efficient conversion of droplet charge to gas-phase ions requires the production of very small and highly charged droplets. Gomez and Tang (23,24) have studied ES- produced droplets with phase Doppler anemometry and flash shadowgraph techniques and were able to deduce the droplet size and charge from measured droplet velocities. On the basis of such findings, Fernandez de la Mora (25)has proposed a relationship for the radius R of the ESproduced droplets from sprays

where p is the density, Vf is the volume flow rate, and y is the surface tension of t h e solvent. Because a

small radius is desirable, increasing the current I by increasing the volume flow rate Vf (see Equation 4 and Figure 2) is counterproductive; Vf also increases the radius of the droplets, as shown in Equation 7. Low flow rates in the range of 1-10 mL/ min, which lead to small droplets (23,24), also lead to optimum ESMS conditions. This is evident by t h e broad maximum of the analyte BH' intensity in the flow range 1-10 mL/ min in Figure 2b. The lack of correlation between the capillary current I and the ion intensity I(BH+) demonstrated in Figure 2 represents a case in which the variable V, leads to a change off; the fraction of droplet charge that is converted to gas-phase charge (i.e., gas-phase ions) and a lack of correlation are expected. Variables t h a t do not change f appreciably lead to good correlations between the capillary current and the observed total gas-phase ion current. As we will show in the final section of this article, the electrolyte concentration is such a variable, but only over a limited range of concentrations. Shrinkage of charged ES droplets The initial charge and size of droplets produced in E S depend on the spray conditions used, as explained in the preceding section. Here we will consider droplets produced a t low flow rates, typically - 5 pL/min at total electrolyte concentrations not exceeding M. These conditions are probably the best for high gas-phase ion yields and high ESMS sensitivity. Furthermore, work by Gomez and Tang (23, 24) and Davis and co-workers (26) h a s provided good d a t a for droplets obtained a t these conditions. The droplets are small and have a narrow distribution of sizes, so they can be considered monodisperse. The size distribution peaks a t radius Ro = 1.5 pm, and t h e droplets have a charge of Qo = C, which corresponds to N = 50,000 singly charged ions (23,

24). The Rayleigh equation (271, which gives t h e condition i n which t h e charge Q becomes just sufficient to overcome the surface tension y that holds the droplet together, is

Q:

Figure 2. Capillary current and analyte ion intensity dependence on solution flow rate. (a) Dependence of capillary current I on solvent flow rate through the capillary. Solid line is cocaine HCI M in methanol; dashed line gives slope predicted by Equation 5. (b) Dependence of massat analyzed ion intensity (countds) of protonated cocaine = BH' on flow rate. Trends are different for current land intensity of BH' with flow rate. (Reprinted from Reference 14.)

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ANALYTICAL CHEMISTRY, VOL. 65, NO. 22,NOVEMBER 15,1993

= 1.25x lO-loR; (8b) where eo is the permittivity of vacuum and y is t h e surface tension. Equation 8b gives the condition for methanol, whose y = 0.0226 N/m2; the numerical factor is for Q in cou-

lombs and R in meters. The initial charge Q , = C observed by Gomez and Tang corresponds to only 50% of the Rayleigh limit charge QR for R, = 1.5 mm. Gomez and Tang (24), who also studied the charge-to-volume ratio of larger ES-produced droplets, found that Q0 comes closer to the Rayleigh limit for larger droplets. Droplets within t h e micrometer range or larger are known to maintain their charge; they do not emit gas-phase ions (24, 26). The droplets s h r i n k by evaporation of solvent molecules until they come close to the Rayleigh limit where they become unstable and undergo fission into smaller droplets (Figure 3). Recent work has shown that droplets with sizes in the l-pm range fission somewhat before, at - 80% of the Rayleigh limit (23, 24, 26). The droplet will shrink by solvent evapo ration at constant Q, until the radius R meets t h a t condition, a n d then a fission will occur. Another very important observation (24, 26) concerns the type of fission that predominates. The droplet does not split evenly into two smaller droplets of approximately equal m a s s a n d charge. Typically, the droplets are observed to vibrate alternately from oblate to prolate shapes. These elastic vibrations stimulate disruptions in which the “parent” droplet emits a tail of much smaller offspring drop lets. An illustration of such a fission is shown in the inset i n Figure 3. This disruption pattern is similar to the disruption at the tip of the Taylor cone. The emitted stream of offspring droplets carries off only about 2% of the mass of the parent droplet but about 15%of the parent’s charge. The radius of the offspring droplets, which a r e quite monodisperse, is roughly one-tenth of the radius of the parent (22,24).A simple calculation (28) shows t h a t a 2% loss of mass leads to - 20 offspring droplets of this radius. The offspring are not only much smaller than the parent but also have a much higher chargeto-mass ratio. I n the subsequent discussion we will call a fission in which such a jet of small offspring droplets is emitted “uneven” fission; “even” fission will mean the formation of two droplets of similar size and charge. The time required for the parent droplet to reach t h e size R , t h a t leads to the first fission can be estimated (28) with use of expressions providing the rate of solvent evapor a t i o n from small droplets (29). When relatively volatile solvents

U.UY

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Figure 3. Schematic representation of time history of parent and offspring droplets. Droplet at top left is a typical parent droplet created near the ES capillary tip at low flow rates. Evaporation of solvent at constant charge leads to uneven fission. The numbers beside the droplets give radius R (pm) and number of elementary charges N on droplet; A t corresponds to the time required for evaporative droplet shrinkage to size where fission occurs. Only the first three successive fissions of a parent droplet are shown. At bottom right, the uneven fission of an offspring droplet to produce offspring droplets is shown. The time scale is based on Equation 9b for methanol evaporation, which produces only a rough estimate. Inset: Tracing of photograph by Gomez and Tang of droplet undergoing “uneven” fission (24). Typical droplet loses 2% of its mass, producing some 20 smaller droplets that carry off 15% of the parent charge. (Inset reprinted with permission of Gomez and Tang.)

such as methanol, acetonitrile, and water are involved and the droplets are a few micrometers in diameter or smaller, the evaporation rate follows t h e surface evaporation limit law (29), which leads to a simple dependence (26, 27) of the droplet radius on time t

R = Ro - 1.2 x

t

(9b) where ij is the average molecular velocity of the solvent gas, Po is the vapor pressure of the solvent at the temperature of the droplet, M is the molar mass of the solvent molecules, p is the density of the solvent, R, is the gas constant, and T i s the temperature of the droplet. The droplet temperature will be about 10 “C lower than the temperature of the ambient gas (air) due t o cooling of the droplet by evaporation (28).The condensation coefficient is a,and its value is - 0.04 for both water and ethanol (28, 29). Equation 9b was obtained by substituting the corresponding parame ters for methanol and assuming that a = 0.04 for methanol also. The predictions of Equation 9b are only intended to provide “ballpark” esti-

mates. Results obtained with Equation 26 from Reference 28 are shown in Figure 3, which gives the time dependence of Q and R for a methanol droplet of initial R, = 1.5 pm and Qo= 8 x C, in ambient air of - 35 “C. The time required for the first fission is - 400 ps. The change of R due to the fission is very small and barely perceptible in Figure 3. Subsequent fissions that occur when t h e p a r e n t droplet again reaches 80% of the Rayleigh limit are included in Figure 3. These fissions require shorter times, - 60 ps, which decrease progressively from fission to fission. A typical offspring droplet has a r a d i u s R - 0.08 pm a n d a n u m b e r of e l e m e n t a r y charges N = 280 at birth. It reaches 80% of the Rayleigh limit in - 40 ps, where its charge is the same and its radius has shrunk to R 0.03 pm. It is the offspring droplets that are expected to ultimately lead to gas-phase ions. The time scale for evolution to gasphase ion-producing droplets indicated by the results in Figure 3 is in the hundreds of microseconds, not very much shorter than the total residence time of the charged droplets in the atmospheric region. The exact residence time depends on the type of ES interface used. It is probably in

ANALYTICAL CHEMISTRY, VOL. 65,NO. 22, NOVEMBER 15, 1993

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REPORT the one to a few milliseconds range for typical interfaces that one would use in ESMS (13). The direct dependence of the rate of evaporation on the vapor pressure of the solvent (Equation 9a) indicates that solvents of relatively low vapor pressure will lead to acceptable ESMS sensitivity only at elevated temperatures. The vapor pressure for methanol used to obtain Equation 9b is 100 torr, which is a usefbl indicator for the required vapor pressure. Unfortunately, because of their small size, the type of fission t h a t offspring droplets undergo has not been observed. The type of fission, even or uneven, is of great interest because it determines the ratio of charges (excess positive ions near the surface) N,relative to n, which is the number of charge-paired ions in the solution inside the droplet. For the case in which even fission occurs (i.e., a droplet separates into two droplets of near equal charge and mass), there will be no increase in Nln, which corresponds to the ratio of the number of excess positive ions near the surface to the number of ion pairs in the bulk. On the other hand, uneven fission such as that discussed above involving 15%of the charge but only 2% of the mass (or volume) in the offspring droplets leads to a n increase of Nln by a factor of 7. This number is arrived at by assuming that when the Offspring! droplets a r e formed the surface charge “slips over” the droplet bulk solution so t h a t the bulk electrolyte concentration in the offspring droplet remains about t h e same as that in the parent drop (28). If uneven fission also occurs for the offspring droplets, Nln for the second generation offspring is further increased by a factor of - 7, leading to a total factor of - 49. The increase of Nln by uneven fission is probably an essential process t h a t makes the production of gasphase ions, and thus ESMS, possible. If all droplet fission were even, a very large electrolyte concentration increase in the droplets could be expected to occur before the droplets reached the small radius R < 10 nm below which gas - phase ion formation is expected to occur. Thus, for the case of droplets t h a t measure R = 1.5 pm and Q = 8 x C considered in Figure 3, a concentration increase by a factor F = 4000 due to the required solvent evaporation can be calculated for the even fission process, leading to droplets that measure R = 10 nm. For more information and for a description of the

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method for evaluating fractions F, see Reference 26. An initial concentration of 5 x lov3 M would lead to a final bulk concentration C ==: 20 M. The formation of droplets with such a high concentration by solvent evaporation at near ambient tempera t u r e is very unlikely because a slowing down of evaporation will occur at high concentrations. Furthermore, even if evaporation did occur, f u r t h e r shrinkage of t h e droplet would lead to charged solid particles from which gas - phase ion emission would not be expected. An analysis of the solute residue problem described above, which leads t o similar conclusions, was first given by Iribarne. See Figure 6 in Reference 30a for additional information. On the other hand, as illustrated i n Figure 3, uneven fission t h a t

forms a first offspring followed by uneven fission of that offspring leads to a droplet of R = 3 nm with N = 2 charges. The calculated concentration increase factor for these droplets is only F = 140 or 30 times smaller than for 10-nm droplets obtained with even fission. Assuming an initial electrolyte concentration of io-* M, one calculates t h e number of charge-paired electrolyte ions n ==: 1 using t h e known volume a n d F = 140. The evaporation of solvent from droplets of this type will lead directly to gas - phase ions! This observation will be elaborated on in the next section, which deals with the mechanism of gas-phase ion formation. I t is useful to restate some of the preceding observations. Even fission treats all droplets equally and, because charge is not enriched, much


solvent evaporation is required. Uneven fission separates the droplets into two classes. The small offspring are enriched in charge and are the future gas-phase ion emitters; the large parent droplets become t h e dumps for t h e unwanted chargepaired electrolyte ions. Nature of the process leading to formation of gas-phase ions Two different mechanisms have been proposed to account for the formation of gas-phase ions from the small charged droplets. The first one depends on the formation of extremely small droplets, R = 1nm, which contain only one ion. Solvent evaporation from such a droplet will lead to conversion of the droplet to a gasphase ion. Such a mechanism was assumed by Dole (31),the first investigator of gas-phase ion production by ES. A more detailed consideration and support of this mechanism was given by Rollgen (32).We shall call this model the single ion in droplet theory (SIDT). This early work did not provide any details on how such droplets should be formed or whether the process should include selectivity that may favor the formation of gasphase ions A+ relative to B’ as a result of specific physicochemical differences between these two species of ions. The other mechanism, proposed by Iribarne and Thomson (301,assumes ion evaporation (emission) from very small and highly charged droplets. Typically, the droplets from which ion emission becomes competitive with Rayleigh fission have a radius of R = 8 nm and N ==: 70 elementary charges (30).Under these conditions, the droplet does not undergo fission but emits gas-phase ions. As N decreases, emission can still be maintained as a result of a decrease of R by solvent evaporation. Thus, t h e Iribarne mechanism does not require the production of very small droplets ( R ==: 1 nm) that contain only one ion. Iribarne emission can occur even when the droplet contains other solutes such as charge-paired electrolytes. This last difference can be used as the basis of experimental tests. In experiments in which a single solute such as NaCl is used, one might expect, particularly a t higher concentrations, t h a t some of the very small SIDT droplets would also contain some charge-paired NaCl species. Solvent evaporation from such droplets should t h e n lead to Na+ (NaCl), ions if SIDT holds. Mass spectra obtained with NaCl solutions

showed the presence of Na' but no Na' (NaCl), were detected. Even ions where n = 1 were not observed for solutions with concentrations of NaCl up to lo-" M (14). The absence of Na+ (NaCl), ions was t a k e n as decisive evidence against the SIDT. However, at the time, even fission was assumed to occur (14), and even fission leads to high solute concentrations in the offspring droplets. As shown in the previous section, uneven fission dramatically increases the ratio Nln of charges N to paired ions. We also showed (Figure 3) that a secondgeneration offspring droplet can have a radius of - 3 nm, carry N = 2 charges, and for an initial electrolyte concentration of [NaCl] = M have only one Na' C1- ion pair (n = 1) inside the droplet. At [NaCl] = lo-" M, one expects N = 3 and n = 10. Further evaporation and uneven fission may lead to a third generation of offspring droplets, and some of these may contain only one Na' and no NaC1. Therefore, mass spectra dominated by the Na' are a possibility. The previous discussion shows that uneven fission provides a mechanism for single ions in droplet formation. Although the observed complete absence of Na' (NaCl), ions in t h e m a s s spectra ( 1 4 ) is not e x plained, the difference between experimental observation and prediction is much less drastic. I n other words, the realization that uneven fission is dominant has provided s t r o n g s u p p o r t for t h e SIDT. Willoughby and co-workers have presented a theoretical model (22) that leads to similar predictions. Details of the lribarne ion evaporation theory Iribarne and Thomson (30) based their theory on a derived equation that provides detailed predictions for t h e r a t e of ion emission from t h e charged droplets. I n particular, it predicts the dependence of the rates on the chemical properties of t h e ions. Observed differences in t h e gas-phase ion intensities IA and IB of ions A+ and B' present at equal concentrations in the sprayed solutions were compared with predictions of the theory by Iribarne and Thomson (30b) a n d subsequently by other workers as well ( 4 , 14, 33, 34). I n general, qualitative agreement between experiment and theory was obtained, lending support to the theory on the assumption that the experimentally observed selectivity

cannot be explained by the SIDT. We will show in the next section of this article that an expanded SIDT is also capable of qualitatively predicting experimentally observed selectivities. To understand the significance of the evidence, one needs to have additional information about t h e Iribarne model and the nature of the experimental evidence. The Iribarne t r e a t m e n t (30) i s

based on transition state theory. The rate constant k , for emission of ions from droplets is kT k, = - exp(-AG'/RT) (10) h where k is the Boltzmann constant, T is the temperature of the droplet, and h is the Planck constant. The free energy of activation, AG', was evaluated on the basis of the model

Figure 4. The "ion evaporation" transition state and its dependence on the ion solvation energy. (a) Transition state for ion evaporation (emission) proposed by lribarne and Thomson. In the initial state, a droplet of radius R has N ionic charges. Typical values are R = 8 nm, N = 70. Distance of ions from surface d is due to solvation shells of ions; see enlarged inset. In the transition state, the escaping ion-solvent molecule cluster M+(SI), is shown at distance x, outside the droplet; x, is typically 0.6 nm. (b) Free energy of escaping ion cluster M+(SI), as a function of distance x obtained with the lribarne equation (30).Maximum of A@ corresponds to the activation energy A@. Values of rate constants k, given in s-' units at T = 298 K droplet temperature. Individual curves are for different solvation energies of the ion solvent molecule cluster M+(SI)m


979 A

REPORT shown in Figure 4a. The transition state selected by the authors resembles the products rather than the initial state (Le., it is a "late" transition state). The advantage of this choice is that the energy of such a state can be expressed with a closed equation based on classical electrostatics and thermodynamics. However, the transition state could be occurring earlier, for example, as the ion disrupts the droplet surface. The energy of such an early transition state would , be much more difficult to evaluate. If a higher free energy barrier did occur at that earlier stage, the predictions of the Iribarne model would be invalidated. The barrier in the Iribarne transition state is due to the opposing electrostatic forces: the repulsion of the escaping ion by the other charges of the droplet and the attraction between the escaping ion and the droplet because of the polarizability of the solvent medium of the droplet. The attraction is larger at short distances between the ion and the droplet surface, but it falls off faster than the repulsion as the distance is increased. The equation for AG* was found (30) to depend on four parameters. The first two are N , the number of charges on the droplet, and R, the radius of the droplet. The rate constant k , increases with N and decreases with R. The other two parameters express the specific properties of the ions involved. The escaping ion in Figure 4a is not the naked electrolyte ion M' but an ion-solvent molecule cluster M'(Sl), containing m solvent molecules. AG' is lowest not for the naked ion M', but when the ion takes several solvent molecules with it. Recall that the transfer of a naked ion such as Na' from aqueous solution to the gas phase (Equation 1) requires a large amount of energy (- 98 kcal/mol). On the other hand, the transfer of Na'(H,O), requires only - 56 kcal/mol. The transfer free energies from gas phase to solution for the naked alkali ions M' and the clusters M' (H,O), are given in Table 11. These data are based on experimental thermochemical measurements available in the literature ( I , 2) and were evaluated in Reference 28. The strongly solvated ions Li' and Na' have large transfer energies (see Table II), and for these ions the Iribarne equation predicts large activation barriers (i.e., low rate constants k , ) . The change of k , with the value of the ion cluster solvation energy is illustrated in Figure 4b, which depicts evaluated barriers 980 A

1.25

1.5

1

Figure 5. I ypical dependence of capillary current I and analyte A' ion intensity in experiments in which a single analyte is added to the solvent. A constant concentration of impurity M Na') present in reagent-grade methanol provides the dominant electrolyte at low analyte concentrations. Results shown are based on morphine H+ as A+. (Reprinted from References 21 and 28.)

for different values of the ion cluster solvation energy. The second parameter t h a t expresses the individuality of the ions is d , the distance of the ion charges from the surface of the droplet (see Figure 4a). The charges cannot be on


t h e droplet surface, where t h e y would minimize their repulsion, because that would partly disrupt their solvation shell. Strongly solvated ions such as Li' hold on strongly to a larger number of solvent molecules and have a larger d. Radii obtained

from ion mobilities in solution (Stokes radii) provide suitable values for d (28).Such values are given in Table 11. Rate constants kI evaluated for a droplet with R = 8 nm with N = 70 charges at a temperature T of 298 K for the alkali ions are given in Table 11. These d a t a illustrate that t h e t r a n s f e r e n e r g y of t h e c l u s t e r M+(H,O), g i v e n b y -AGosol [M’(H,O),] in Table I1 is the decisive parameter for the value of kI. Li’, which h a s the largest cluster transfer energy, has the lowest k,. The k, for Cs’ is larger by a factor of - 5000 relative to k , for Li’. These results predict that large differences in the ESMS sensitivities of these ions are expected provided that the Iribarne theory is valid. Iribarne and Thomson (30b)have presented calculations of the relative k , for the alkali cations. Their results (Table 11) show a much s m a l l e r spread of kI values from Li’ to Cs’ and are much closer to the experimental results. However, the calculated values (28) are probably more trustworthy because they were obtained with experimental data for AG”,,, [M’(H,O),] and d. Iribarne and Thomson used a classical electrostatic expression of questionable validity for the solvation energy and estimated the values for d. Also given in Table I1 are values for the experimental sensitivity coefficients k obtained from the experim e n t . T h e significance of t h e s e values and how they can be compared with predictions of the Iribarne theory and t h e SIDT are described in the next section. Dependence of ion intensities on concentration In this section we will examine how the analyte ion intensity determined with the mass spectrometer depends on the concentration of the analyte ion in the solution t h a t is electrosprayed and how this intensity is affected by the presence of other electrolytes. O t h e r electrolytes a r e almost always present as impurities in the solvent used, as other coanalytes, and as buffers required in reversed-phase LC or CE. The question concerning the expected ion intensity for a given analyte under given concentration conditions is part of the daily practice of ESMS. Trying to understand the factors that determine the observed intensity leads one rapidly to ask questions about t h e mechanism of ES discussed in previous sections. Experimentally determined ion in-

tensities at different concentrations of analytes in t h e solution t h a t is electrosprayed are shown in Figures 5-7. The intensity of a single analyte observed in mass spectra obtained with a series of solutions with increasing analyte concentration is shown in Figure 5. The logarithmic plot u s e d accommodates a wide range of ion intensities and concentrations. The analyte plot has a typical shape that is commonly observed (14,28, 35). A linear section with a slope of = 1 in the low concentration range up to M is followed by “saturation” and a small decrease of intensity at the highest concentration M). The linear section, where the intensity IA of the analyte A is proportional to the concentration [A’], is a region useful for the quantitation of A+ in the solution. Such linearity is generally observed (14,28, 35, 36).

The key to understanding the complete analyte curve is to realize that the solution involved is not a single electrolyte system because of the solvent that is used. Unless special deionization procedures are imple mented, impurity electrolytes are always present. The impurities in the reagent-grade methanol used (Figures 5-7) are primarily ammonium and sodium salts at a total concentration of - lop5 M. For analyte concentrations below M, ES is possible only because of the presence of impurity electrolyte. In the low concentration range of analyte A, the capillary current I is carried by t h e dominant electrolyte-the impurities B, which are at a constant concentration. I is therefore constant in this range (see Figure 5a). The sum of the mass analyzed total ion intensitites Itot= IA + IB is also constant in this range be-

Figure 6. Results from a series of experiments in which a given analyte A is at M, and NH,CI is increased at concentrations constant concentration, [A+] = shown. (a) Change of capillary current Iwith increase of NH4CI concentration. Essentially the same Iis observed with any of the analytes A. (b) Change of ion intensity of mass analyzed A+ corrected for mass dependent transmission of A+. Only one A+ and NH4CI were present in a given series in which [NH4CI]was increased. Solid lines are curves predicted by Equation 12. Analytes A+ are identified above and in Table 111. Best fit ratios k,lk, used with Equation 12 are given in Table Ill. (Adapted from Reference 28.)


981 A

REPOR7 cause it is dominated by the impurities B. Above an analyte A concentration of - low5M, the analyte begins to dominate and the total electrolyte concentration begins to increase. An increase of capillary current I occurs in this region as expected from Equation 5.

I= ((,C)" n = 0.2-0.3 (11) Because the molar conductivity hk for different electrolytes M' gener ally changes by less than a factor of 2 over a wide range of electrolytes (37) and the exponent n is very small, the current I is essentially independent of t h e n a t u r e of t h e electrolyte. Changes of I with concentration C must be considered because C can be changed over a much wider range. However, again as a result of the small exponent n, these changes are still relatively small (a factor of 2-4) for a change of C by a factor of 100. The shape of the Itot curve is seen to be very similar to that of the capillary current I (Figure 5 ) . An approximate proportionality between the two currents is generally observed in this concentration range (14,223). For an analyte concentration above lo-' M, the intensity IB of the impurity ions B' is observed to decrease (Figure 5). This decrease is a consequence of t h e weak dependence of the current I on the total concentration. Because the current is proportional to the total droplet charge, the addition of A barely increases t h e droplet charge. However, i n t h e droplets, the ions A+ compete with the ions B' in the conversion process to gas-phase ions. A proportionality to concentrations of ions A+ and B' in the droplets may be expected in this competition so that an increase of [A+] should lead to a decrease of production of gas-phase ions B', that is to a decrease of IB. This is exactly what is observed in Figure 5. On t h e basis of such considerations, Equation 12a has been proposed (20, 26) for a two-electrolyte system; Equation 12b is for a threeelectrolyte system.

The equations for IB a n d I , are analogous to those for IA. For simplicity we will discuss t h e two982 A

.

electrolyte equation. IA+ is the mass spectrometrically detected ion intensity of ion A+ corrected for the mass dependent transmission of the mass analyzer. [A'] and [B'I are the electrolyte concentrations in the solution to be electrosprayed. The product f p is a factor that was assumed to be independent of the chemical nature of t h e ions; f is t h e fraction of charges on the droplets that are converted to gas-phase ions and # is the ion- sampling efficiency or the fraction of t h e ions detected with t h e mass spectrometer relative to t h e g a s - p h a s e ions produced by t h e droplets at 1 atm. Values of f p can be determined with Equation 13, IA+ IB = f p I (13) which follows from Equation 12a, by measuring the total mass analyzed ion current, IA + IB, and the capillary current I .

In experiments in which concentrations and the nature of the electrolytes were changed, the product f p was found to be approximately constant (28)up to a total concentration of - 5 x M. A rough estimate of f = 0.3 has been obtained (28) for a 20-pL/m flow rate with methanol as the solvent. A decrease of f p is generally observed above - 5 x lop4 M (Figure 5 ) . It has been attributed (14)to a decrease off, which is a result of an expected increase in droplet size at high electrolyte concentrations (19).The value of p depends on the design of the interface between the atmospheric region and the vac-


uum of the mass analyzer; values of = 10-4-10-5 may be expected (28). I t should be noted that according to Equation 12a, the intensity IA depends only on the ratio kA/kB and not on the individual values of kA and kB. The expression kA/kB is the yield ratio of gas-phase ions A+ and B' relative to the solution concentrations. In earlier work such ratios were generally obtained (2, 30, 33) by determining the ion intensities IA and IB with solutions where [A'] = [B']. These measurements can be considered as a special case in the use of Equation 12, which when applied to IA and IB leads to

p

A '

IB

=

kA

-for[A]=[B] k*

(14)

The use of Equation 12 permits the determinations of kA/kB for a wider set of concentration conditions where [AI # [Bl. The dependence of IA and IB on the solution concentrations [A"] and [B'] is expected to be indirect. Much more relevant would be t h e number N A and N B of charges due to ions A' and B+ at the surface of the droplets and the number nA and HB of ions A' and B+ that a r e paired with negative counterions in the droplets. These numbers are not available. However, Equation 12 can be used for a second purpose. The extent of agreement or deviation of Equation 12 with experiments in which the concentrations [A] and [Bl are changed over wide limits permits deductions to be made on the relationship between the concentrations [AI and [Bl and the relative numbers of ions in the droplets, Recently, Smith and co - workers (35)have shown that an equation of the same form as Equation 12 can be usefully applied to ES involving effluent from CE. Experiments in which only a single electrolyte A is added to the solvent that contains impurity B, as in the case for the data used in Figure 5 , can be successfully fitted by Equation 13 (see Reference 20) but do not provide an accurate determination of kA/kB and a sensitive indication of whether t h e value kA/kB remains constant with changes of concentration of A. Results where two electrolytes are added to the solvent Better tests are obtained when two electrolytes, A and B, are added to the solvent. The presence of electro-

lyte impurities in the solvent, which will be called C, makes this a system of three electrolytes. Results from such experiments (28) are given in Figure 6. In these experiments, the concentration of a given analyte A was kept constant, [A+] = 1 x M, and the concentration of B', which is always NH",, was M. The increased above 2 x solid curves in Figure 6, which provided a good fit to the experimental points (28), were obtained w i t h Equation 12. The kA/kB values obtained (28)from these fits were used to acquire some of the data shown in Table 111. The plots obtained in Figure 6 illustrate a situation of practical importance, that is, how the intensity of an analyte A+ is suppressed by the presence of a second electrolyte B, which is used as a buffer. Substantial decreases of I A result. When A+ is Bu,N+, the ion with the highest kA, the loss of intensity is by a factor of 3; this is the smallest decrease obtained. When Cs+, which has t h e lowest kA, is present, intensity decreases by a factor of 12 when [NH",] is increased from to M. Equation 12 predicts that the decrease of I A with [B+]will be greater as the value of kB increases. NH", has a low value of k , and therefore the decreases of IA observed in Figure 6 are not the worst case. When the ion with the highest coefficient, Bu,N+, was used instead of NH",, t h e observed decrease of the analyte intensities for Cs" and MorH' was much higher, by a factor of - 200 (28). Therefore, in the actual practice of ES, buffers that have cations with low coefficients k should be used if possible.

More extensive experiments, particularly under conditions in which t h e added electrolytes a r e a t t h e same concentration [A"] = [B'], have shown (28) that in general a fit with Equations 12 a n d 14 cannot be obtained over the complete concentration range 10-8-10-3 M with a single value for the ratio kA/kB. The results i n Figure 7 illustrate t h e three typical cases observed. The ratio kA/kB remains constant over t h e complete concentration range when kA = kB. This is the case for K" and Cs+ (Figure 7a). Note that for the logarithmic plots used, a constant kA/kB corresponds to a constant vertical distance between log I A and log IB. The joint decrease of I A and IB at low [A] = [B] is due to the presence of electrolyte impurity in the solvent and is analogous to the decrease discussed for Figure 5. When kA/kB > 1 and approximately constant at a high concentration, a lower kA/kB is observed at low concentrations. kA/kB = 1 is observed for t h e tetraalkylammonium cations Pen4N+ and Et,N+ in the low concentration range (Figure 7b), whereas for the pair and n-C,H,,NH,+ and Cs+, a decrease to a lower and approximately constant value at low concentration is observed, but this value is > 1 (Figure 74. The factors responsible for t h e above changes have been explored (28) in some detail only from t h e standpoint of the Iribarne theory, which was extended to include the effect of surface activity. The high kA/kB observed at high concentrations are due to A+ having a higher Iribarne rate constant and/or higher surface activity. The effect of differential surface activity is to increase

the ratio of surface charges NA/NB above the ratio expected on the basis of the concentrations [AI/ [Bl. The rate of ion emission from an Iribarne droplet is proportional to the number of surface charges and will be higher for the more surface-active ion even when the Iribarne rate constants are equal. The observed decrease of kA/kB at low concentrations was attributed to depletion of the ion A+ (28), which has a higher evaporation rate. At high concentrations, the ions in the solution of the droplet dominate over the surface charges (i.e., n >> N ) . The faster ion evaporation of A+ from t h e surface does not decrease t h e surface charge ratio NA/NB because A+ ions are rapidly supplied from the droplet bulk where nA/nB = [A]/[Bl. At low concentrations, t h e r e a r e many fewer ions i n t h e bulk, although the surface charge number is approximately the same; this leads to the condition N > n. Faster evaporation of A+ from the surface leads to a relative depletion of A+ over B+ in the droplet and the faster emission of A+ cannot be maintained. Therefore, the observed kA/kB decreases at low concentrations. The situation illustrated i n Figure 7b may occur when t h e difference between t h e Iribarne rate constants for A+ and B+ takes precedence over the difference in surface activities. For the case in Figure 7c, the reverse was assumed. This argument was based on the rationale that the relative number of surface charges NA/NB adjusts to the relative surface activities already in the parent droplets (see Figure 3), where t h e number of bulk ions is large so that this process will be less affected by depletion. Comparison of coefficients with lribarne theory and SlDT Comparisons of the experimental coefficients with the Iribarne rate constants and expected surface activities should be made only in the high concentration range, where depletion is absent. The experimental coefficients given in Table I11 are averaged results from experiments shown in Figure 6 and experiments in which [A] = [Bl, as in Figure 7. The range of coefficient values in Table I11 is not very large. From the standpoint of ESMS this is a desired result and means that we can expect to detect with a fair sensitivity any analyte ion present in the solution. The ions with the highest coefficients i n Table I11 a r e t h e large tetraalkyl ammonium ions, which are expected to have low solvation


983 A

REPORT energies and thus high Iribarne rate constants. They also have relatively high surface activities (38).Actually, a general correlation between low solvation energies and high surface activities is expected. Because of the absence of suitable quantitative information on solvation energies and surface activities in the literature, it is difficult to establish the extent to which each factor contributes to the value of the experimental coefficient (Table 111). I t appears that ions with no surface activity such as the alkali ions are best suited for comparison with theory. As discussed earlier, t h e agreement between the calculated kI and the experimental values for the alkali ions is not very good (see Table 11). The qualitative agreement observed between the Iribarne pre dictions and experiments involving organic ions ( 4 , 14, 31, 32) might be due to surface differences only, and if this is the case, these results need not be considered as support of the Iribarne theory. The SIDT, as originally stated (31, 32), did not provide criteria for selectivity on the basis of physicochemical properties of the ions. I t is obvious that the surface activity is the factor that should lead to selectivity for the SIDT. Ions that are enriched on the surface will preferentially end up with the final droplet that contains a single ion. An analysis examining how well the depletion phenomena can be rationalized on the basis of the SIDT, expanded to include surface activity, has not been made yet, but some depletion could be expected. At present, it is not possible to state with certainty which theory, ion evaporation or SIDT, fits better with the available evidence. However, a shift in favor of the SIDT has occurred. Fortunately, from a practical standpoint, qualitative predictions for high selectivity on the basis of low solvation energy or high surface activity will both be valid, because the two parameters are positively correlated. Emission of gas-phase ions from the Taylor tip of the ES capillary Very recently Siu and co-workers (39, 40) have proposed that the gasphase ions are not formed from the charged droplets but a r e emitted from the Taylor cone tip. This is a radical proposal that, if it turns out to be true, would require a n extensive reinterpretation of experimental observations described in the preceding discussion. The evidence presented (39, 40) is 984 A

indirect. D e t e r m i n a t i o n s of t h e ESMS spectra of cytochrome c and myoglobin in 0.2% acetic acid and in 0.2% propionic acid solutions in wat e r showed t h a t t h e ion intensity profiles resulting from multiple pro tonation were almost identical when acetic or propionic acid was used (39). The authors argued t h a t the profiles obtained with the two acids should be different if the ions originated from droplets t h a t had been subjected to extensive evaporation, They estimated t h a t the hydrogen ion concentration in the acetic acid solution should increase because evaporation increases the acetic acid content in water. For propionic acid, a decrease of acidity was estimated. With these assumptions one can expect that the polyprotonated ion profile in the ESMS spectra should be shifted to higher proton numbers for the acetic acid solution spectra. Be-

Figure 7. Observed mass analyzed ion intensities /A and /B for experiments in which solution concentrations of analytes [A+] = [B'] were increased. Typical cases discussed in text are illustrated in (a)-(c). A decrease of IA and le intensities observed at [A'] = [B'] < lop5 M is a result of competition from electrolyte impurity C ' in methanol at constant concentration [C'] 2: M. (Adapted from Reference 28.)


cause this was not the case, the ions must originate from solution t h a t has experienced minimal evaporation. The most likely source for such ions is the Taylor cone (39). However, the relative vaporization estimates were made with the assumption that liquid-vapor equilib rium data apply to evaporation from very small droplets of a volatile liquid such as water. This assumption is not justified. The rate of vaporization from such droplets is not thermodynamically controlled-it is kinetically controlled (29).The socalled surface - controlled evaporation holds. The composition of the vapor escaping under surface control from the very small droplets of water and 0.2% acetic or propionic acid cannot be predicted with equilibrium data, and therefore the evidence (39) for the absence of evaporation is weak. E S mass spectra obtained with ethylene glycol and water as solvents were compared (40).When the liquid in the ES capillary was at room temperature, water was found to lead to normal ion intensities; those observed with the much less volatile ethylene glycol were 200 times lower. In experiments in which the spray probe was at 100 "C, the intensities with ethylene glycol as the solvent were found to be 40 times lower than those observed with water. Higher intensities from water relative to ethylene glycol are expected on the basis of ion production from droplets that have to reduce their size by evaporation. However, the authors provide evidence, largely based on computer modeling, that the intensity differences expected between water a n d ethylene glycol would have been much larger had the ion production depended on solvent evaporation from droplets, because the expected evaporation of ethylene glycol is insignificant. Even if the arguments (40)are accepted at face value, they do not prove that the ion emission for water is not from droplets that have experienced evaporation. They only show that the observed very weak ion production from glycerol may be from the Taylor cone. Ion emission may in general be occurring from the Taylor cone and from droplets. For volatile solvents, ion production from t h e charged droplets should dominate over Taylor cone emission. Nonvolatile solvents such as glycerol may lead only to the weak Taylor cone emission. Ion emission from liquid tips in a vacuum is known to occur and is the basis of electrodynamic MS ( 4 1 ) . The ion currents obtained

a r e several orders of magnitude smaller than the capillary current I in ES. The work of Siu and co-workers is interesting and provocative; however, its significance will have to be established in the future. Mechanisms for formation of multiply-charged macroions The preceding part of this article applies to conventional ions, that is, singly- and doubly-charged inorganic or organic ions of molecular mass up to a few hundred Daltons. Although many ESMS applications in chemistry and biochemistry involve such ions, the ability of ES to produce multiply -charged macroions such as the polyprotonated peptides and proteins is also an area of great interest (4-8). The mechanistic aspects of this area deserve a separate article. Here, we can consider briefly only two aspects that are related to the previous discussion. For simplic ity we consider only polyprotonated peptides and proteins. We side with the view (5)that the ESMS observed degree of protonation is related to the number of basic groups that are protonated in the solution. However, the degree of protonation will be modified by changes of pH due to solvent and acid evaporation from the charged droplets (42, 43). As pointed out in the previous section, it is not easy to establish how the pH will change on evaporation from small droplets, particularly when volatile acids are used. The Iribarne theory for ion evaporation from a charged droplet can be applied to multiply - charged macro ions. A model has been suggested by Fenn (44). In the initial state, the polyprotonated macromolecule is lo cated inside the small charged droplet. All charges of the macromolecule are neutralized by negative counter ions of t h e solution. Because of Brownian motion, a given part of the macromolecule can approach t h e droplet surface. As this happens one or more protonated sites of the macromolecule will replace charges on the surface of the droplet. Thermal activation may provide the energy for one of t h e charged sites of the macromolecule to move some distance ouside the droplet. This process is similar to the activated step of the Iribarne theory involving the normal ion mechanism. Once such a charge is past the activation barrier, the separation between it and the charges on the droplets leads to repulsion, which facilitates the escape of a portion of the molecule that car-

ries a second charged site. The two charges then facilitate the escape of other charged parts and ultimately the escape of the whole macroion. An alternative mechanism can be proposed on the basis of the observed uneven fission and presence of oscillations ( 2 4 ) i n t h e small charged droplets. For example, one of the offspring droplets in Figure 3 had a radius of 30 nm and some 280 charges. Assume that a polyprotonated protein with - 30 protonated sites is present in the droplet. Some of the protonated sites may be near t h e surface of t h e droplet where they would not be charge balanced by

negative counterions; thus, they can be part of the 280 surface charges. An elastic oscillation of the droplet may expose additional protonated sites, precipitating a special uneven fission in which the polyprotonated species with only a small amount of solvent becomes the small and highly charged offspring of the uneven fission. The distinction between this model and that of Fenn-Iribarne is that in Fenn’s model the activation energy is thermal, and in the present model it is the presence of elastic deformations t h a t “activates” t h e event. References (1) (a) Kebarle, P. Annu. Rev. Phys. Chemistry 1977,28,445; (b) Keesee, R. G.; Castleman, A. W. J. Phys. Chem. Re5 Data 1986, 15, 1011; (c) Chandrasekhar, J.; Spellmeyer, D. C.; Jorgensen, W. L. J. Am. Chem. SOC.1984, 106,903. (2) Desnoyers, J. E.; Joliceur, C. In Modern Aspects of Electrochemistry; Bockris, J.O.M.; Conway, B. E., Eds.; Plenum Press: New York 1969; Vol. 5, p. 20. (3) (a) Yamashita, M.; Fenn, J. B. Phys. Chem. 1984, 88, 4451; (b) Yamashita, M.; Fenn, J. B. Phys. Chem. 1984, 88, 4671. (4) Fenn, J. B.; Mann, M.; Meng, C. K.;

Wong, S. F.; Whitehouse, C. M. Science 1989,246, 64. (5) Covey, T. R.; Bonner, R. F.; Shushan, B. I.; Henion, J. Rapid Commun. Mass Spectrom. 1988,2, 249. (6) Smith, R. D.; Olivares, J. A.; Nguyen, N. T.; Udseth, H. R. Anal. Chem. 1988, 60, 436. (7) (a) Mirza, A. U.; Cohen, S. L.; Chait, B. T. Anal. Chem. 1993,65,1; (b) Winger, B. E.; Light-Wahl, K. J.; Rockwood, A. L.; Smith, R. D. J. A m , Chem. SOC. 1992, 114,5897. (8) Beu, S. C.; Senko, M. W.; Quinn, J. P.; Wampler, F. M.; McLafferty, F. W. J. Am. SOC. Mass Spectrom. 1993,4,557. (9) Bajley, A. G. Electrostattc Spraying of Liauzds: J o h n Wilev a n d Sons: New Ygrk, 1988. (10) Loeb, L. B.; Kip, A. F.; Hudson, G. G.; Bennett, W. H. Phys. Rev. 1941,60, 714. (ll)Pfeifer, R. J.; Hendricks, C. D. AL4.A J 1968,6,496. (12) Taylor, G. I. Proc. R. SOC.London A 1964, A280,383. (13) Ikonomou, M. G.; Blades, A. T.; Kebarle, P. Anal. Chem. 1990, 62,957. (14) Ikonomou, M. G.; Blades, A. T.; Kebarle, P. Anal. Chem. 1991, 63, 1989. (15) Ikonomou, M. G.; Blades, A. T.; Kebarle, P. J Am. SOC.Mass Spectrom. 1991, 2, 497. (16) Wampler, F. W.; Blades, A. T.; Kebarle, P. J. Am. SOC.Mass Spectrom. 1993, 4, 289. (17) Blades, A. T.; Ikonomou, M. G.; Kebarle, P. Anal. Chem. 1991, 63, 2109. (18) Van Berkel, G. J.; McLuckey, S. A,; Glish, G. L. Anal. Chem. 1992, 64, 1586. (19) Hayati, I.; Bailey, A. I.; Tadros, T. F. J. Colloid Inte$ace Sci. 1987, 117,205 and 222. (20) Smith, D.P.H. IEEE Truns. Ind. Appl. 1986, LA-22, 527. (21) Tang, L.; Kebarle, P. Anal. Chem. 1991, 63,. 2709. (22) (a) Willoughby, R.; Sheehan, E.; Jarrell, A.; Masecic, T.; Pedder, R.; Penn, S. Presented at the Annual Meeting of the American Society for Mass Spectrometry, San Francisco, CA, 1993; (b) Sheehan, E. W.; Willoughby, R. Presented at the Annual Meeting of the American Society for Mass Spectrometry, S a n Francisco, CA, 1993. (23) Gomez, A.; Tang, K. Proceedircgs of the Fifth International Confirence on Liquid Atomization and Spray Systems I (CLASS-91); Smerjians, H. G., Ed.; NIST: Gaithersburg, MD; Special publication 813. (24) Gomez, A.; Tang, K. Physics of Fluids, in press. (25) Fernandez de la Mora, J.; RosselLiompart, J. Proceedings of the 39th ASMS Conference on Mass Spectrometry and Allied Topics; ASMS: Nashville, TN, 1991; p. 441. (26) (a) Taflin, D. C.; Ward, T. L.; Davis, E. J. Langmuir 1989, 5, 376; (b) Davis, E. J . ISA Trans. 1987,26, 1. (27) Lord Rayleigh Philos. Mag. 1882, 14, 184. (28) Tang, L.; Kebarle, P., submitted for publication in Anal. Chem. (29) Davies, C. N. In Fundamentals ofAerosol Science; Shaw, D. T., Ed.; John Wiley and Sons: New York, 1978; p. 154. (30) (a) Iribarne, J. V.; Thomson, B. A. J. Chem. Phys. 1976, 64, 2287; (b) Thomson, B. A.; Iribarne, J. V. J. Chem. Phys. 1979, 71, 4451. (31) Dole, M.; Mack, L. L.; Hines, R. L.; Mobley, R. C.; Ferguson, L. D.; Alice, M. B. J Chem. Phys. 1968,49,2240. (32) (a) Rollgen, F. W.; Bramer-Wegner,


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REPORT E.; Buttering, L. J. Phys. Colloq. 1984,45, Supplement 12, C9-297; (b) Schmelzeisen-Redeker, G.; Buttering, L; Rollgen, F. W. Int. J. Mass Spectrom. Ion Processes 1989,90, 139. (33) (a) Sakairi, M.; Yergey, A. L.; Siu, K.W.M.; LeBlanc, J.C.Y.; Guevremont, R., Berman, R. S. Anal. Sci. 1991,7,199; (b) Rafaelli, A.; Bruins, A. P. Rapid Commun. Mass Spectrom. 1991,5,269. (34) Hiraoka, K. Rapid Commun. Mass Spectrom. 1992,6,463. (35) Smith, R. D.; Wahl, J. H.; Goodlett, D. R.; Hofstadler, J. A. Anal. Chem. 1993,65,574 A. (36) Rafaeli, A.; Bruins, A. P. Rapid Commun. Mass Spectrum. 1991,5,269. (37) Landolt, B. Zahlenwerle und Functionen; Springer Verlag: Berlin, 1960; Vol. 2, pp. 366, 533, 651. (38) Tamaki, K. Bull. Chem. SOC. Jpn. 1967, 40, 38. (39) Siu, K.W.M.; Guevremont, R.; LeBlanc, J.C.Y.; O’Brien, R. T.; Berman, S. S. Org. Mass Spectrom. 1993,28, 579. (40) Guevremont, R.; LeBlanc, J.C.Y.; Siu, K.W.M.Org. Mass Spectrom., i n press. (41) Cook, K. D. Mass Spectrom. Rev. 1986, 5, 467. (42) Kelly, M. A.; Vestling, M. M.; Fenselau, C. C.; Smith, P. B. Org. Mass Spectrom. 1992,27,1143. (43) Ashton, D. S.; Bedell, C. R.; Cooper, D. J.; Green, B. N.; Oliver, R.W.A. Org. Mass Spectrom. 1993,28,721. (44) Fenn, J. B. J. Am. Soc. Mass Spectrom. 1993,4, 524.

tained by any other means. He is the recipient of the Medal of the Canadian Chemical Society (1986) and the ACS Frank H. Field and Joe L. Franklin Award for Outstanding Achievement in Mass Spectrometry (1994). I

I

Paul Kebarle received a Dipl. Eng. Chemistry (1952) fiom ETH Zurich (Switzerland) and his Ph.D. in physical chemistry (1956)f’rom the University of British Columbia (Canada). After postdoctoral work at NRC in Ottawa, he joined the Department of Chemistry at the University of Alberta. His early research in ion-molecule reactions at near-atmospheric pressures led to the discovery that ion-molecule equilibria can be determined by MS techniques. The ion thermochemistry resulting f’rom equilibrium measurements has become a main source of data on the energetics of ions in the gas phase. These data find application in many areas, including ions in solution. His interest in ESMS was based on the recognition that it affords the production of many ion species in the gas phase that cannot be ob-

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Liang Tang was born in Shanghai (China) and received both his B.Sc. degree in chemistry and his M.Sc. degree in environmental engineering at East China University of Chemical Technology, Recently he received his Ph.D. in analytical chemistry at the University of Alberta, where he worked under the direction of Paul Kebarle. His research interests range from environmental chemistry to MS, with emphasis on analytical techniques including chemical ionization, atmospheric pressure ionization, fast atom bombardment, and electrospray.

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From ions in solution to ions in the gas phase - the mechanism of

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