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. 2023 Jun 12;127(24):5445–5452. doi: 10.1021/acs.jpcb.3c02002

Local pH at Nonionic and Zwitterionic Lipid/Water Interfaces Revealed by Heterodyne-Detected Electronic Sum-Frequency Generation: A Unified View to Predict Interfacial pH of Biomembranes

Achintya Kundu †, Shoichi Yamaguchi †,§,*, Tahei Tahara †,‡,*
PMCID: PMC10292198  PMID: 37308160

Abstract

graphic file with name jp3c02002_0007.jpg

For biomembranes, which are composed of neutral as well as charged lipids, the local pH at lipid/water interfaces is extremely important in their structural formation and functional activity. In our previous study of the charged lipid/water interfaces, we found that the local pH at the interface is governed by the positive or negative sign of the charge of the lipid: i.e., the local pH is dictated by the repulsive or attractive electrostatic interaction between the charged lipid headgroup and the proton. Because of the lack of net charge in the headgroup of the neutral lipid, the factor determining the local pH at neutral lipid/water interfaces is less straightforward, and therefore it is more challenging to predict the local pH. Here we apply heterodyne-detected electronic sum frequency generation (HD-ESFG) spectroscopy to nonionic and zwitterionic lipids to investigate the local pH at the neutral lipid/water interfaces. The obtained results indicate that the local pH at the nonionic lipid/water interface is higher than in bulk water by 0.8 whereas the local pH at the zwitterionic lipid/water interface is lower by 0.6, although the latter is subject to significant uncertainty. The present HD-ESFG study on neutral lipids, combined with the previous study on charged lipids, presents a unified view to consider the local pH at biomembranes based on the balance between the electrostatic interaction and the hydrophobicity provided by the lipid.

Introduction

Biological membranes not only exclude undesirable substances migrating from the extracellular region but also selectively transport necessary agents such as protons, ions, and so on into the cell. Therefore, the interface region of biological membranes plays critical roles in various biochemical processes that are essential for life.1,2 For the part that the charged lipids form, their charged headgroup creates an electric field, resulting in the surface potential3 that is the difference in electrostatic potential between bulk water and the charged interface. Depending on the surface charges, the proton is either attracted or repelled from the interface so that the pH at the interface is shifted from that in bulk.4,5 The local pH and the surface potential at the membrane interface are important for a variety of biological phenomena such as the binding of proteins and drugs onto membranes,6−8 insertion and orientation of membrane proteins,9−12 and aggregation of proteins and peptides at membranes.13 For example, alamethicin has been extensively studied as a model for large channel proteins, of which the orientation at biological membranes is associated with antimicrobial activity and ion channel forming ability, depending on the local pH.10,12

Because of such relevance to biochemical reactions, there have been numerous attempts to experimentally and theoretically evaluate the local pH and surface potential.3,13−20 A number of experimental approaches were used to estimate the local pH and the surface potential at aqueous interfaces, including nonlinear optical spectroscopy,19,22−26 the vibrating plate capacitor method,27−29 X-ray photoelectron spectroscopy,30 nuclear magnetic resonance,31 conductance measurements,32 atomic force microscopy,33 voltage sensitive fluorescent probes,34 and electron paramagnetic resonance.35 Among them, second harmonic generation (SHG) and sum frequency generation (SFG) have received much attention for estimating the local pH and the surface potential.26,36−42 In particular, we developed heterodyne-detected electronic sum frequency generation (HD-ESFG) spectroscopy43,44 that allows one to measure interface-selective electronic spectra for quantitative pH spectrometry.4,23,24 We applied HD-ESFG to charged lipid/water interfaces in our previous study, demonstrating that the local pH at the charged lipid/water interfaces is governed by the positive or negative sign of the headgroup charge through the electrostatic interaction.

In the present study, we investigate the local pH at neutral lipid/water interfaces by using HD-ESFG. The interface pH spectrometry allows for concluding that the local pH at the nonionic lipid/water interface is higher than in bulk water, whereas that at the zwitterionic lipid/water interface looks lower. We attribute the higher pH at the nonionic lipid/water interface predominantly to a lower local dielectric constant that is not favorable for the proton. The lower pH at the zwitterionic lipid/water interface is considered to result from dipolar electrostatic field that outweighs the effect of lower dielectric constant. Combining these results with our previous study on charged lipids, we provide a unified viewpoint for predicting the local pH near various biomembranes.

Experimental Section

4-Heptadecyl-7-hydroxycoumarin (HHC) was purchased from Fluka and used as received. 1,2-Dipalmitoyl-sn-glycerol (DPG) and a zwitterionic 1,2-dipalmitoyl-sn-glycerol-3-phosphocholine (DPPC) were purchased as powders from Avanti Polar Lipids. Chloroform (99.7%, GC 149 grade) was purchased from Kanto Chemical Co. and used to prepare the lipid solutions. Sodium hydroxide was purchased from Wako and used as received. We used ultrapure water (Millipore, 18.2 MΩ cm resistivity) for the aqueous subphase. The surface pressure was measured with a commercial surface tension meter (Kibron, Inc., Helsinki, Finland). The interface density of the lipids was approximately 2 molecules nm–2, and the lipid monolayers were in the liquid condensed phase. The bulk pH was regulated by the concentration of sodium hydroxide in bulk water and was measured using a commercial pH meter (Horiba, B-212). All the experiments were performed at 296 ± 2 K.

The experimental details of HD-ESFG spectroscopy were described previously.23,24,45 In brief, a Ti:sapphire regenerative amplifier system (Spectra-Physics, Spitfire Pro XP, 120 fs, 1 kHz, 3.5 mJ) was used as the light source. One portion of the amplifier output was used as a narrow-band ω1 pulse (795 nm), and the other was focused into water to generate a broad-band ω2 pulse (540 nm–1.2 μm). The ω1 and ω2 pulses were noncollinearly focused onto the same spot at the lipid/water interface. When the ω1 and ω2 pulses were temporally and spatially overlapped, the sum frequency (ω1 + ω2) was generated at the lipid/water interface. The ω1, ω2, and ω1 + ω2 pulses were again focused by a spherical concave mirror onto a GaAs (110) surface to generate the sum frequency of ω1 and ω2 once more, which acted as a local oscillator field. A fused-silica glass plate with 1 mm thickness was placed between the sample and the concave mirror. The glass substrate delayed the ω1 + ω2 pulse relative to the reflected ω1 pulse by 170 fs due to the frequency-dependent difference in the group velocity. This delay resulted in the time difference between the ω1 + ω2 pulse generated from the sample and that from the GaAs surface. The ω1 + ω2 pulse from the sample and that from GaAs propagated collinearly and sequentially through an analyzing polarizer and entered a polychromator. After being spectrally dispersed, the ω1 + ω2 sum frequency light was detected by a multichannel detector (Roper Scientific, Spec-10:2KBUV), where interference fringes between the two ω1 + ω2 pulses were superimposed on the detected spectra. Fourier analysis of the interference fringes yielded the complex χ(2) (second-order nonlinear optical susceptibility) spectra of the pH indicator at the lipid/water interfaces. The ω1 + ω2, ω1, and ω2 pulses were s-, p-, and s-polarized, respectively.

Results and Discussion

Figure 1a shows the acid–base equilibrium of HHC, which is used as a surface active pH indicator in this study. For neutral lipids to examine, we used two lipids having different head groups with the same acyl chain, i.e., nonionic DPG and a zwitterionic DPPC, shown in Figure 1b. The pH indicator is coadsorbed at the lipid/water interfaces with its alkyl chain aligned with the lipid acyl chains and its chromophore facing interfacial water.4,24

Figure 1.

Figure 1

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(a) Acid–base equilibrium of the surface-active pH indicator (4-heptadecyl-7-hydroxycoumarin, HHC). (b) Chemical structures of the nonionic lipid (1,2-dipalmitoyl-sn-glycerol, DPG) and the zwitterionic lipid (1,2-dipalmitoyl-sn-glycerol-3-phosphocholine, DPPC).

Figure 2 shows the interface-selective electronic χ(2) spectra of the pH indicator at the nonionic DPG/water interface. The imaginary (Im χ(2)) and real (Re χ(2)) parts show absorptive and dispersive band shapes, respectively. The horizontal axis represents the sum frequency (ω1 + ω2) wavelength. Under the present two-photon resonant and one-photon nonresonant condition, the Im χ(2) spectra plotted against the sum frequency wavelength can be interpreted in the same manner as UV–visible absorption spectra.43 The sign of Im χ(2) indicates the absolute orientation of the pH indicator at the interface.45 The Im χ(2) spectrum at pH 12.7 exhibits a negative band peaked at 372 nm, which is due to the conjugate base (A–) of the pH indicator. With lowering bulk pH from 12.7 to 6.2, the intensity of the A– band at 372 nm decreases, while the Im χ(2) signal of the acid (HA) becomes predominant which is seen in the wavelength region shorter than 350 nm. The intensity increase of the HA signal can be more readily seen for Re χ(2) with a higher signal-to-noise ratio (S/N). All the Re χ(2) spectra at different bulk pH pass through an isosbestic point at 380 nm, and simultaneously all the Im χ(2) spectra at different bulk pH seem to pass through an isosbestic point around 350 nm. The presence of these isosbestic points indicates that the χ(2) spectra of the pH indicator are solely due to the acid and the conjugate base of the pH indicator at the DPG/water interface with no change in their total concentration, meaning that we can neglect the adsorption–desorption equilibrium of the pH indicator between at the surface and in the bulk.4,23,24

Figure 2.

Figure 2

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(a) Imaginary and (b) real parts of the electronic χ(2) spectra of the pH indicator HHC at the nonionic lipid DPG/water interface. The black, pink, green, blue, and red lines represent spectra obtained at bulk pH 12.7, 10.9, 9.5, 9.1, and 6.2, respectively. The dashed lines stand for global fits. (See the text for details.) The interface density of the lipids was approximately 2 molecules nm–2, which was 10 times higher than that of HHC.

Figure 3 shows the interface-selective electronic χ(2) spectra of the pH indicator at the zwitterionic DPPC/water interface. We observe the negative A– band peaked at 376 nm in the Im χ(2) spectrum at bulk pH 12.6. With lowering bulk pH from 12.6 to 8.4, the signal due to A– decreases, and in turn, the signal due to HA increases in the wavelength region shorter than 350 nm. Also in this case, the intensity increase of the HA signal is seen more clearly in Re χ(2) with a higher S/N. The Im χ(2) and Re χ(2) spectra at the DPPC/water interface exhibit the isosbestic points at around 350 and 380 nm, respectively, similarly to the DPG/water interface.

Figure 3.

Figure 3

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(a) Imaginary and (b) real parts of the electronic χ(2) spectra of the pH indicator HHC at the zwitterionic lipid DPPC/water interface. The black, pink, green, blue, and red lines represent spectra obtained at bulk pH 12.6, 11.5, 11.3, 10.8, and 8.4, respectively. The dashed lines stand for global fits. (See the text for details.) The interface density of the lipids was approximately 2 molecules nm–2, which was 10 times higher than that of HHC.

The electronic χ(2) spectra at the neutral lipid/water interfaces were analyzed in the same manner as in our previous study on the charged lipid/water interfaces,4 first for the nonionic DPG/water interface. We consider that the χ(2) spectra of the pH indicator shown in Figure 2 are expressed as the linear combination of the HA and A– spectra. The mole fraction of HA, fHA(pH), and that of A–, fA(pH), are a function of bulk pH and can be expressed as follows:

graphic file with name jp3c02002_m001.jpg 1
graphic file with name jp3c02002_m002.jpg 2

Here pH means bulk pH, and x is a parameter (to be determined in a fitting analysis) representing bulk pH at which fHA(pH) and fA(pH) become equal. Each experimental spectrum in Figure 2 is represented by Si(λ), where i (=1, 2, ..., N) is an index of the spectrum at each pH and λ is the wavelength. Then Si(λ) for any i can be expressed as follows:

graphic file with name jp3c02002_m003.jpg 3

where pHi stands for bulk pH at which Si(λ) is measured. Although the spectrum of HA, SHA(λ), and the spectrum of A–, SA(λ), are unknown, they can be expressed by the linear combination of S1(λ) and SN(λ) in the following manner:

graphic file with name jp3c02002_m004.jpg 4
graphic file with name jp3c02002_m005.jpg 5

where cHA1, cHA, cA1, and cA are unknown coefficients of the linear combination (to be determined in a fitting analysis). We fit the χ(2) spectra in Figure 2 using eq 3 (in combination with eqs 1, 2, 4, and 5) as a fitting function by treating x, cHA1, cHA, cA1, and cA as global fitting parameters. Figure 2 demonstrates that all the spectra of the pH indicator at the nonionic DPG/water interface are well reproduced with the global fits (dashed lines). The converged fitting parameters allow us to obtain the pH dependence of the dissociation degree (i.e., fA(pH)) of the pH indicator at the DPG/water interface as shown in Figure 4a, where [HA] and [A–] are equal at bulk pH of 9.3. We note that this pH value, 9.3, cannot be regarded as the true pKa (negative logarithm of the acid dissociation constant) of the pH indicator at the DPG/water interface because the dissociation degree at the interface is plotted against bulk pH in Figure 4a. (Thus, this value is sometimes called “apparent pKa”.23) Thinking that the local pH at the DPG/water interface is shifted from bulk pH by Δ, we can draw the top horizontal axis representing the local pH at the interface in Figure 4a. Then the true pKa of the pH indicator at the DPG/water interface is expressed as 9.3 + Δ.

Figure 4.

Figure 4

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pH dependence of the dissociation degree of the pH indicator (a) at the nonionic DPG/water interface and (b) at the zwitterionic DPPC/water interface. The solid circles represent the dissociation degree determined from the analyses of the χ(2) spectra, and the solid lines represent the fits using eq 2.

We can carry out the same global fitting analysis for the χ(2) spectra of the zwitterionic DPPC/water interface. As shown in Figure 3, the fits well reproduce the experimental data also in this case. Figure 4b shows the pH dependence of the dissociation degree of the pH indicator at the DPPC/water interface. With Δ′ defined as the pH difference between at the DPPC/water interface and in the bulk, the pKa of the pH indicator at the DPPC/water interface is expressed as 11.3 + Δ′.

For the determination of Δ and Δ′, we need to estimate pKa of the pH indicator at the interfaces, for which we can take advantage of the solvatochromism of the A– band as we did in our previous study.4,24 The solvatochromic shift of the A– band was investigated in various solvents by Drummond and Grieser,46 and Figure 5b plots the peak wavelength (λmax) of the A– band against the solvent dielectric constant (ε) according to their pioneering work. In harmony with the negative solvatochromism of HHC, the absorption band of the A– species of HHC exhibits a blue-shift with increasing ε (i.e., solvent polarity). Using this plot, we can estimate the local ε of the interfaces from the peak wavelength of the A– band in the interface-selective electronic spectra shown in Figures 2 and 3. By plotting those peak wavelengths onto the λmax–ε curve in Figure 5b, we estimate that the local ε of the DPG/water interface is 25–1+2 and that of the DPPC/water interface is 20–10, which are in good agreement with reported values.47,48 Then, by using a relation between pKa and ε reported by Fernández and Fromherz,49 we can estimate pKa of the pH indicator at the interface. Figure 5a shows the pKa–ε curve of HHC, where pKa becomes lower with increasing ε because A– is more stabilized in a polar medium than HA. Using the local ε values of the interfaces, we estimate the pKa values of the pH indicator at the DPG/water and DPPC/water interfaces at 10.1–0.2+0.1 and 10.7–0.5, respectively, which are plotted in Figure 5a.

Figure 5.

Figure 5

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(a) Relation between pKa of the pH indicator (HHC) and ε of the surrounding medium. The black circles and the solid line are taken from the upper dashed curve in Figure 3 of ref (49). The black circles represent data points obtained from dioxane–water mixtures of different ratios. The black circle at ε = 78.5 corresponds to neat water where pKa is 7.75. The ε values of the nonionic lipid DPG/water and zwitterionic lipid DPPC/water interfaces give red and green points on the solid line, respectively, which allows us to estimate pKa of the pH indicator at the DPG/water and DPPC/water interfaces as 10.1 and 10.7, respectively. (b) Relation between the peak wavelength of the Im χ(2) spectrum of A– and ε of the surrounding medium. The open circles with error bars and the solid line are taken from Figure 4 of ref (46). Dotted lines placed at ±2 nm from the solid line fully cover the error bars of the individual points, representing an estimate of the experimental error. The open circles represent the data points obtained from different bulk solvents. The peak wavelengths of A– at the DPG/water interface (372 nm) and DPPC/water interface (376 nm) give red and green points on the solid line, respectively, which allows us to estimate the local ε of the DPG/water and DPPC/water interfaces as 25 and 20, respectively.

We can now consider that pKa estimated from the solvatochromism of the pH indicator (Figure 5a) is equal to pKa expressed with the use of Δ (or Δ′) through the pH spectrometry at the interfaces (Figure 4). Then, we obtain Δ = +0.8–0.2+0.1 for the nonionic DPG/water interface, indicating that the local pH is higher than the bulk pH by 0.8. On the other hand, we obtain Δ′ = −0.6–0.5 at the zwitterionic DPPC/water interface. The error of Δ′ is large compared with Δ because the peak wavelength of HHC in this range is not very sensitive to the change of the dielectric constant in the region of ε = 10–20. This low sensitivity causes the large error of the dielectric constant, propagating to the uncertainty of pKa of DPPC. This Δ′ value looks to indicate that the local pH at the zwitterionic DPPC/water interface is lower than bulk pH by 0.6 although this value contains a substantial error.

The local pH represents the local activity of the proton and thus the chemical potential of the proton at the interface. With the electrostatic potential at the interface represented by ϕI, the chemical potential of the proton at the interface, μH+I, is expressed as follows:

graphic file with name jp3c02002_m006.jpg 6

where μH+◦I is the standard chemical potential of the proton at the interface, F is the Faraday constant, R is the gas constant, T is temperature, and aH+ is the activity of the proton at the interface. By replacing “I” with “B”, the chemical potential of the proton in the bulk, μH+B, is expressed in the following manner:

graphic file with name jp3c02002_m007.jpg 7

where μH+◦B is the standard chemical potential of the proton in the bulk and aH+ is the activity of the proton in the bulk. Note that the electrostatic potential in the bulk is zero. Under the adsorption–desorption equilibrium of the proton at the interface and in the bulk (i.e., μH+I = μH+), we can express the pH difference (Δ) between at the interface and in the bulk as follows:

graphic file with name jp3c02002_m008.jpg 8

where δH+ is the difference between the standard chemical potential of the proton at the interface and that in the bulk. This δH+ represents the energy required to transfer the proton from the bulk to the interface, apart from the electrostatic interaction. On the assumption that the local ε at the interface governs δH+, we can use the ε dependence of the standard chemical potential of the proton, which Fernández and Fromherz experimentally determined.49 They measured the mean degenerate activity coefficient of HCl as a function of ε in water/dioxane mixtures, which allows us to draw the δH+–ε curve shown in Figure 6. This curve indicates that δH+ increases as ε decreases. This tendency is readily understood because the proton is less stable in a medium of lower ε. Using the estimated values of the local ε at the interfaces (Figure 5b), we can estimate that δH+/(RT ln 10) at the DPG/water and DPPC/water interfaces are 0.9 and 1.2, respectively, as plotted in Figure 6. Consequently, from eq 8, FϕI/(RT ln 10) at the DPG/water and DPPC/water interfaces are estimated to be −0.1 and −1.8, respectively. This estimation indicates that the DPG/water interface does not exhibit a noticeable electrostatic potential whereas that at the DPPC/water interfaces is significantly negative, although the headgroup of either lipid does not have any net charge. We note that the negative electrostatic potential at the DPPC/water interface is consistent with the overall positive sign of the hydrogen-bonded OH stretch band at a zwitterionic lipid/water interface, which was observed in our previous HD-VSFG study50 and also a molecular dynamics (MD) simulation study,51 although direct hydrogen bonding between the water molecule and the phosphate group of the DPPC may also play a substantial role.

Figure 6.

Figure 6

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ε dependence of δH+ of the proton in water/dioxane mixtures. The black circles represent experimental data points obtained from a study by Fernández and Fromherz.49 The solid line stands for a phenomenological exponential fit. The ε values at the DPG/water and DPPC/water interfaces are shown with red and green points on the solid line, respectively, which allows us to estimate δH+/(RT ln 10) of the proton at the DPG/water and DPPC/water interfaces as 0.9 and 1.2, respectively.

Equation 8 separates the origin of the pH difference between at the interface and in the bulk into two parts. One is the difference between the standard chemical potential of the proton at the interface and that in the bulk (i.e., δH+), which is commonly called the medium effect.49 The other is the electrostatic potential at the interface (i.e., ϕI). As estimated above, the pH difference between at the nonionic DPG/water interface and in the bulk is mostly due to the medium effect with minimal contribution from the electrostatic potential. This result is compatible with our intuition: the nonionic interface just prepares a nonpolar environment with negligible electrostatic potential, resulting in a lower concentration of the proton at the interface than in the bulk.

Compared to nonionic DPG, it is more challenging to understand why the local pH at the zwitterionic DPPC/water interface looks lower than bulk pH, although the upper limit of the experimental error might make the local pH higher. The estimated values of δH+ and ϕI at the DPPC/water interface clearly shows that the most important difference between the DPPC/water and DPG/water interfaces is the electrostatic potential. We may attribute the negative value of ϕI at the DPPC/water interface to the charge distribution in the headgroup: the average position of the nitrogen atom in the cationic choline group along the interface normal is closer to bulk water than that of the phosphorus atom in the anionic phosphate group, resulting in the formation of an electric double layer to create the negative electrostatic potential at the interface compared to the potential at the bulk water phase. Actually an MD simulation study demonstrated that zwitterionic and anionic lipids exhibit almost the same negative electrostatic potential at the interface.17 Although we need to leave the quantitative discussion of the electrostatic potential at the zwitterionic DPPC/water interface for a future study where more elaborate experiments will minimize the error bar, we can qualitatively argue that the dipolar nature of the DPPC headgroup is the key to understand the difference between the local pH at the DPG/water and DPPC/water interfaces.

Table 1 summarizes the local pH and related properties of the nonionic and zwitterionic lipid/water interfaces evaluated in the present study, along with those at the cationic and anionic lipid/water interfaces studied in our previous work using the same pH indicator.4Table 1 clearly demonstrates that the ionic property of lipids is the dominant factor for the local pH through the electrostatic potential at the interface: the electrostatic potential at the nonionic DPG/water interface is negligible as expected, whereas that at the zwitterionic DPPC/water interface is negative as much as the negatively charged DPPG/water interface. Consistently, the electrostatic potential at the positively charged DPTAP/water interface is positive with almost the same absolute value as the negatively charged DPPG/water interface. The medium effect on the local pH is always positive ranging from +0.5 to +1.7. This suggests that the nonpolar nature of the headgroup solvating the pH indicator is common to these lipids regardless of their ionic properties, which has an effect to repel the proton from the interface and make interfacial pH higher. Table 1 provides a useful data set not only for predicting the local pH, ε, and electrostatic potential at biomembranes but also for carrying out MD simulation of the lipid/water interfaces with targeting molecular-level understanding of various phenomena occurring at the biomembranes.

Table 1. Local pH, Local Relative Dielectric Constant (ε), Local Standard Chemical Potential of Proton (δH+), and Electrostatic Potential (ϕI) at Nonionic, Zwitterionic, Cationic, and Anionic Lipid/Water Interfaces.

lipid local pH (interface pH – bulk pH) local relative dielectric constant (ε) local standard chemical potential of proton (δH+/(RT ln 10)) electrostatic potential (FϕI/(RT ln 10))
DPG (nonionic) +0.8 25 +0.9 –0.1
DPPC (zwitterionic) –0.6 20 +1.2 –1.8
DPTAP (cationic)a +3.8 13 +1.7 +2.1
DPPG (anionic)a –1.5 35 +0.5 –2.0
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a

The data of DPTAP and DPPG are taken from our previous paper.4

Conclusions

We studied the local pH at the aqueous interfaces of nonionic DPG and zwitterionic DPPC as representative neutral lipid interfaces. We carried out pH spectrometry at the interfaces with HD-ESFG and measured pH dependence of the dissociation degree of a pH indicator. The obtained results enabled us to estimate pKa of the pH indicator at the interface and pH difference from the bulk that was contained as a parameter in the expression of the acid–base equilibrium of the pH indicator. Briefly, based on the solvatochromic shift of the electronic transition of the pH indicator, we estimated the effective relative dielectric constant ε at the interface and obtained pKa of the pH indicator at the interface using the ε–pKa relationship reported in a bulk study. The comparison between this “true” pKa value with the “apparent” pKa value determined by the pH spectrometry using HD-ESFG made it possible to estimate the local pH at the DPG/water and DPPC/water interfaces: The local pH at the nonionic DPG/water interface is higher than bulk pH by 0.8 whereas that at the zwitterionic DPPC/water interface is lower than bulk pH by 0.6. The local pH at the DPG/water interface can be rationalized simply by the medium effect that has a tendency to repel the proton from the interface. The local pH that looks lower than the bulk pH at the DPPC/water interface is attributable to the negative electrostatic potential originating from the zwitterionic charge distribution in the headgroup. The present study of the nonionic and zwitterionic lipids, in combination with our previous work on the anionic and cationic lipids, provides a basis for considering and predicting the local pH at a biomembrane based on its lipid composition.

Acknowledgments

This work was supported by a JSPS KAKENHI (Grants 25104005, 24245006, 18H05265, 25288014, 18H05265, and 23H00292).

The authors declare no competing financial interest.

Special Issue

Published as part of The Journal of Physical Chemistry virtual special issue “Hiro-o Hamaguchi Festschrift”.

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