Useful Information

ILSB Method

Here you can find information on the method used to determine Gibbs transfer energy values of single ions necessary for the $\mathrm{pH_{abs}^{H_{2}O}}$ and $\mathrm{\mathit{E}_{abs}^{H_{2}O}}$ scales.

The method is based on measurements of the electrical potential difference of two half-cells connected by a salt bridge that is filled with a pure ionic liquid (ILSB). As example, cell I shows the case of the redox system Ag⁺/Ag, a schem of which is given below.

Ag │AgZ^m (c, S₂) │ [N₂₂₂₅][NTf₂] │ AgZⁿ (yc, S₁)│Ag                                                        cell I

Some video turtorials on how to conduct the measurements can be found here.

The cell potential of cell I reads as:

With x is the sum of both LJPs at the ends of the ILSB. The cation and the anion of the used ionic liquid amyl­triethylammonium bis(trifluoromethanesulfonyl)imide [N₂₂₂₅][NTf₂] show in the neat IL identical and in the investigated solvents S nearly identical diffusion coefficients, enabling that a major part of the liquid junction potentials (LJPs) at the phase boundaries becomes negligible in good approximation. With the assumption that the Gibbs transfer energy of the [N₂₂₂₅]⁺ and the [NTf₂]⁻ ion between the considered solvents is identical the potential difference of cell I is related to the Gibbs transfer ernergy of the Ag⁺ ion according to Eq. m2.

An assessment of above assumption can be found here.

Two necessary criteria have to be fullfiled when dealing with single ion transfer energies: (i) The addition of cation and anion values have to give the value of the resulting electrolyte, which is an observable as neutral chemical entity. (ii) Consistency of the values, meaning that one value (e.g. of the transfer from water to acetonitrile, $\Delta_{\mathrm{tr}}G{^\circ}\left(\mathrm{Ag}^{+},\mathrm{H_{2}O\rightarrow MeCN}\right)$) have to turn out as the sum of two other (e.g. of the transfer from water to propylene carbonate and propylene carbonate to acetonitrile, $\Delta_{\mathrm{tr}}G{^\circ}\left(\mathrm{Ag}^{+},\mathrm{H_{2}O\rightarrow PC}\right)$ + $\Delta_{\mathrm{tr}}G{^\circ}\left(\mathrm{Ag}^{+},\mathrm{PC\rightarrow MeCN}\right)$).

Criterion (i) is assessed here.

Criterion (ii) is an inherent part of the ILSB method: Different implementations of cell I varying in S₁ and S₂ lead to a network of measurements (see video below) that can be analyzed with the method of least squares. We express the consistency as $\sigma$ (as it is essentially the rms) as part of the result of this analysis; the lower $\sigma$ the higher the consistency, which also indicates the precision of the measurements, including the assessement whether the approximations used are justified. The ratio m/n between measured values m and unknowns n, that is, the number of half-cells, of a network is a criterion for the reliability of the results. The higher this ration the more robust is the network. You also can find more information on the network analysis here.

Note: Eq. m1 contains not the standard Gibbs transfer energy. However, at low electrolyte concentrations and calculation of (relative) activity coefficients (by the Debye-Hückel Theory or the MSA (Mean Spherical Approximation)) standard values can be obtained.

Note further, that the redox system also can be (and concerning criterion (i) have to be at least for an anionic system) varied, either in both half-cells or in one half-cell only. A network can be constructed that is increased in m/n. Since this is ongoing work the values might change a little on occasion.

Comparison with some Extra-Thermodynamic Assumptions

Here you can find an (rather heuristic) assessment of the assumption on which the ILSB method is based. This also includes a comparison with other extra-thermodynamic assumptions. In this context, also criterion (i) is examined.

The assumption of the ILSB method is the equality of the Gibbs transfer energies of the cation and anion of the IL used in the salt-bridge:

The better Eq. a1 is achieved, the smaller are the electric potential contributions of the LJPs at both ends of cell I to the potential of cell I. If Eq. a1 holds exactly, the LJPs vanishes completely, and the assumption is equivalent to the negligible LJP assumption (n-LJP) of Parker et al.^[1] The measured cell potential E_I can be compared to the cell potentials measured by Parker et al., cell P.

Ag │AgClO₄ (0.01 mol L⁻¹, S₂) │ [NEt₄][Pic] (0.1 mol L⁻¹, S₁ or S₂) │ AgClO₄ (0.01 mol L⁻¹, S₁)│Ag                cell P

The Table below shows the differences for the available redox systems and solvent pairs.

In cases where the transfer occurs from a protic to an aprotic solvent (or vice versa), the agreement is less than in the other cases; it does not fall within the (estimated) uncertainty of the n-LJP assumption (excepting the transfer from MeOH to PC). Parker et al. estimate a greater degree of uncertainty (unquantified) for cases involving water than for the other cases (approximately 20 mV). Details can be found here.

On the other hand, the TATB assumptions is given by Eq. a2.

[TA]⁺ is the tetraphenylarsonium cation and [TB]⁻ is the tetraphenylborate anion. One instantly can identify the equality to our assumption. If both assumptions are valid, the coefficients a and b in Eq. a3 would be 0 and 1, respectively.

a and b can be obtained by plotting the Gibbs transfer values obtained with the TATB method, $^\mathrm{TATB}\Delta_{\mathrm{tr}}G°\left(\mathrm{i,S_{1}\rightarrow S_{2}}\right)$,^[2] vs. that obtained with the ILSB method, $^\mathrm{ILSB}\Delta_{\mathrm{tr}}G°\left(\mathrm{i,S_{1}\rightarrow S_{2}}\right)$. The Table below shows the results for the Gibbs transfer energies of the single ions i = Ag⁺, Cl⁻. S₁ and S₂ are H₂O, MeCN, DMF, PC, EtOH, and MeOH. R is the correlation coefficient of the linear regression.

Thus, the Gibbs transfer energies obtained with both assumption can be considered as being equal within the given uncertainties of the methods. The separation of the data between S₁ = H₂O and S₁ ≠ H₂O shows that the reported asymmetrical hydration of the [TA]⁺and the [TB]⁻ ions of about 5 kJ mol^{−1 [3]} can be sensed with the ILSB method. Details can be found here.

Next, the sum of the Gibbs transfer energies of a single cation and anion must give the Gibbs transfer energy of the resulting electrolyte. The latter can be obtained without extra-thermodynamic assumptions. In case of Ag⁺ and Cl⁻ the Table below shows the outcome of this juxtaposition. Additionally, self determined values with cell III as well as values taken from literature of the transfer of AgCl, derived from pK_sp values are listed.^[1,4]

Ag │AgZ (c, S) │ [N₂₂₂₅][NTf₂] │ CiCl (c, S)│AgCl │Ag              cell III

All values are given as $\Delta_{\mathrm{tr}}G{^\circ}\left(\mathrm{i},\mathrm{H_{2}O\rightarrow S}\right)$ in kJ mol⁻¹.

The ILSB values are consistent (and from the outcome above also the TATB values), and in accordance with literature data. Thus, the approximations applied to the ILSB method are reasonable. Details can be found here.

A similar outcome is obtained when data concerning proton transfer values are analyzed. Details can be found here.

Concerning the assumptions underlying to the ILSB method and the TATB method, both also can be false, however, in this case both must be biased by the same magnitude – a more complex scenario. According to Ockham´s Razor, we therefore prefer the scenario in which both assumptions are correct (within the given uncertainty).

Literature

[1] R. Alexander, A. J. Parker, J. H. Sharp, W. E. Waghorne, J. Am. Chem. Soc. 1972, 94, 1148.

[2] Y. Marcus, Pure Appl. Chem. 1983, 55, 977.

[3] R. Scheu, B. M. Rankin, Y. Chen, K. C. Jena, D. Ben-Amotz, S. Roke, Angew. Chem. Int. Ed. 2014, 53, 9560.

[4] M. K. Chantooni, I. M. Kolthoff, J. Phys. Chem. 1973, 77, 1.

Terminology

The terminology we use differs in part from the one conventionally used in redox chemistry. Thus, it needs to be evaluated if a new terminology is needed or if it is unnecessary and confusing.

From a didactical standpoint, the historically evolved terminology within redox chemistry is far from being intuitive compared with the terminology used in acid base chemistry. Here, the focus on the proton as being the reason for an “acid” or “acidity” resulting in the expressions “protonation” and “deprotonation” is self-explanatory.

On the contrary, the quality “reducing” or “reductive” often is seen as being inherent to the molecule, which is able to “reduce” another one; from this phrase it remains unclear that an electron is transferred in this process. Furthermore, in redox chemistry the counterpart to this expression is “oxidizing”. Likewise, it is not evident that an electron is the actual principal player in an oxidation process. The fact that a reducing agent is oxidized concomitantly with an oxidizing agent being reduced whereby the direction of electron transfer is not intuitive during a redox reaction, creates confusion for beginners in chemistry, for non-chemists and even for professionals.

We therefore adopted the terminology of Bockris and Reddy^[1] they used in their textbook of electrochemistry in the 1970s and also found entry into the IUPAC gold book^[2] “electronation” and “deelectronation” being as self-explaining as “protonation” and “deprotonation”. To complete the analogy to acid/base chemistry, which is also covered in this article, we use the expression “reducity” to indicate the reducing power of a medium.

This was done for three reasons: 1) conciseness, 2) accentuation of the analogy to acid–base chemistry, and 3) distinction from the general term “reduction”, which also translates to “decrease” or “diminishment of something” and therefore is misleading—except for those who remember the historic background of the reaction of metals with oxygen, the oxidation, and the return, the reduction, of metal oxides to metals.

Below you can find a brief summary of the expressions we use.

We are quite aware that introducing new expressions may lead to even more confusion than to more clarity. However, we are convinced that in redox chemistry more clarity concerning nomenclature is required. This can be achieved, if the parallelism to the acid–base chemistry is utilized by focusing on the particle of interest: the view on the proton with its acidity has to be given a counterpart. This is the focus on the electron with its “reducity”. Likewise, the “oxidity” is opposed to the basicity, but these terms should be avoided, since both are in fact redundant and lead to confusion.

Literature

[1] J. O. M. Bockris and A. K. N. Reddy, Modern Electrochemistry, Plenum Press, New York, 1970.

[2] M. Nič, J. Jirát, B. Košata, A. Jenkins, A. McNaught (Eds.) IUPAC Compendium of Chemical Terminology, IUPAC, Research Triagle Park, NC, 2009.

Standard States

The standard states used here are listed in the following. They are in accordance with IUPAC recommendations.^[1]

T = 298.15 K

p = 10⁵ Pa

All values are given in the mol L⁻¹ scale.

Gibbs solvation energies are given as the transfer of particles from p = 10⁵ Pa in the gaseous state to c = 1 mol L⁻¹ in the solvation state.

We use the EC-B convention according to Bartmess.^[2]

We use as main anchor point the value for the hydration of the proton as determined by Tissandier et al.:^[3]

Δ_solvG°(H⁺, H₂O) = −1104.5 kJ mol⁻¹.

However, the value recommended by Hünenberger et al. is quite similar:^[4]

 Δ_solvG°(H⁺, H₂O) = −1100.0 kJ mol⁻¹.

Note that this value is given in the EC-FD convention and in the mol kg⁻¹ scale. For the proton, the EC-B value is shifted by about −3.6 kJ mol⁻¹ with respect to the EC-FD value (for other ions the shift is negligible). For hydration, the distinction between the mol L⁻¹ and the mol kg⁻¹ scale is considered to be negligible (ca. 0.007 kJ mol⁻¹). Changing the mol L⁻¹ scale into the mol kg⁻¹ scale, the pH_abs value changes by 1 pH unit for the tenfold increase or decrease of the solvents density compared to 1 kg L⁻¹. Most solvents lie in the range of 0.75 – 1.25 kg L⁻¹, thus the changes lie in the range of about −0.1 to +0.1 pH units depending on the solvent.

For values taken from the literature, we cannot know with certainty those states that are not specified exactly. In such cases, we have attempted to deduce the standard states from the context. It should be noted, however, that values published before 1994 cannot of course be given in the EC-FD convention. To the best of our knowledge and belief, the Δ_solvG°-  and Δ_trG°-values listed on this website all refer to the same standard states.

The detailed description can be found under

V. Radtke et al., Pure Appl. Chem. 2021, 93(9), 1049.

Literature

[1] M. B. Ewing, T. H. Lilley, G. M. Olofsson, M. T. Ratzsch, G. Somsen, Pure Appl. Chem. 1994, 66, 533.

[2] J. E. Bartmess, J. Phys. Chem. 1994, 98, 6420.

[3] M. D. Tissandier, K. A. Cowen, W. Y. Feng, E. Gundlach, M. H. Cohen, A. D. Earhart, J. V. Coe, T. R. Tuttle, J. Phys. Chem. A 1998, 102, 7787.

[4] P. Hünenberger, M. Reif, Single-Ion Solvation. Experimental and Theoretical Approaches to Elusive Thermodynamic Quantities, The Royal Society of Chemistry, Cambridge, 2011.

$\boldsymbol{\mu}$ or $\boldsymbol{\alpha}$

Here we briefly outline the reasons for and the implications of selecting $\mathit{\mu}$ or $\mathit{\alpha}$ as the basis of the definition of the pH and pe scales. Some parts are taken from own publications.

The (intrinsic) chemical potential $\mathit{\mu}$ of a chemical entity is a fundamental thermodynamic state function that reflects its chemical environment (i.e. its bonding conditions, etc.), and allows conclusions to be drawn about how this entity will react to changes in this environment.^[1] In two of the most important reaction classes in chemistry, namely acid-base and redox reactions, ions, i.e. electrically charged entities, are involved. Because of their charge, these entities are affected by electric fields in addition to their chemical environment, and the corresponding state function characterizing these entities is the electrochemical potenbtial $\widetilde{\mu}$ ,^[2] commonly expressed as the sum of $\mathit{\mu}$ and $\mathit{\varphi}$ , the inner or Galvani potential of the phase in which the entity is located. $\mathit{\varphi}$ is the sum of the outer or Volta potential $\mathit{\psi}$ and the surface potential $\mathit{\chi}$ of this phase. The latter is caused e.g. by dipole orientations at the phase´s surface. The real potential $\mathit{\alpha}$ , defined as the sum of $\mathit{\mu}$ and $\mathit{\chi}$ , is contrasted here with $\mathit{\mu}$.

When dealing with electrically charged entities some aspects must be considered that are absent for electrically neutral entities or electrically neutral pairs of entities (e.g. electrolytes), for which $\mathit{\mu}$ are observables. First and foremost, it is impossible to prepare and analyze single ion solutions due to the macroscopic electroneutrality condition.^[3] Similarly, the definition of the chemical potential of a charged entity is thermodynamically not straightforward. For these reasons, the definition of $\mathrm{pH_{S}}=-\lg a_{\mathrm{S}}\left(\mathrm{H^{+},S}\right)$ referring to the activity $\textit{a}$ (which is a $\mathit{\mu}$ -based value) of the proton in S is accepted as notional. With the same reasoning the Nernst equation, containing activities of the redox active species, and thus, the standard hydrogen electrode (SHE_S) as zero point of the redox potential scale, must also be regarded as a notional definition of the electrode potential. As long as S is unchanged, this approach works reasonably well. In case of pH measurements, IUPAC recommends a measuring protocol for S = H₂O that takes into account the effect of counterions. However, if S is made different and values are to be related to each other, IUPAC prefers $\mathit{\alpha}$ , but explicitly only for the $\mathrm{SHE_{H_{2}O}}$ (goldbook entry: standard hydrogen electrode, absolute electrode potential, a corresponding entry for the absolute pH value is not available).^[4] The reason is the measurability of $\mathit{\alpha}$ in contrast to $\mathit{\mu}$ . For consistency, however, both scales should be defined on the same quantity. Thus, if the pH was defined with $\mathit{\alpha}$ instead of $\mathit{\mu}$ , it would be accepted that the surface potential influences the acidity of a medium, changing it by about 3 pH units in the case of the aqueous phase. It is therefore worth analyzing $\mathit{\alpha}$ in more detail: For the definition of $\mathrm{SHE_{H_{2}O}}$ the standard real potential $\mathit{\alpha}°$ was used, i.e.

This formula may implicitly contain the assumption that the surface potential $\mathit{\chi}$ is independent of the composition of the solution. To the best of our knowledge, there have been no reports on whether and how the surface potential changes with the composition of a phase. However, we expect this to change, as, for example, in water, the presence and type of ions can be structure-breaking or structure-forming.^[5] This obviously also includes the structure of the surface and thus the surface potential. In this case, Eq. $\mathit{\mu}1$ would have to be revised:

With $\mathit{\chi}°$ being the surface potential of the pure phase adjactent to the vacuum at 298.15 K. Since $\mathit{\mu}$ changes according to

with the composition of the phase, i.e. with the activity of the entity, we can expect something similar for $\mathit{\chi}$ , in the simplest scenario:

$\mathit{\chi(a)}$ is the surface potential as a function of the activity (and thus of the chemical potential of the entity!), which has yet to be determined. (We do not expect $\chi=\chi{^\circ}+RT\ln a$. In contrary, it seems to be more reasonable to expect a dependence on the ionic strength, i.e. $\mathit{\chi(\mathrm{I})}$ , or, even more complicated, one has to consider $\mathit{\chi}\left(a_{\mathrm{e}},a_{\mathrm{f}},a_{\mathrm{g}},\ldots\right)$ where $\textit{a}_\mathrm{e}$ is the activity of the chemical entity under consideration, and $\textit{a}_\mathrm{f}$, $\textit{a}_\mathrm{g}$, …, are the activities of other (charged) chemical entities in the phase.) In other words: The above-mentioned 3 pH units are valid under standard conditions; for any particular non-standard condition, the value will differ individually from 3 pH units. As long as $\mathit{\chi(a)}$ is unknown (or $\mathit{\chi}$ has not yet been determined for the particular non-standard condition, or it has not been shown that $\mathit{\chi(a)}=0$ ), such a pH or pe scale is of little use. If one takes the position (or even shows) that $\mathit{\chi}$ is constant, i.e., does not depend on the activity, then one might as well take the position that $\mathit{\chi}=0$ , since adding a constant value to a pH or pe value (which must first be determined at great expense) is unnecessary. The same applies in principle to Gibbs transfer energies: if different but constant $\mathit{\chi}$ values are assumed for different phases, two constant values are added instead of one, but this could also be omitted.

In this context, the following question needs to be answered: Does the surface potential play a role when, for example, in the case of acidity, it is a subject of discussion to what extent protons catalyze a reaction? Or, to put it another way: Are homogeneous chemical reactions in solution influenced by electrostatic fields? If the answer is “yes,” it is worthwhile to determine both $\mathit{\chi}°$ and $\mathit{\chi(a)}$ (for as many solvents as possible). If the answer is “no,” it is sufficient to know $\mathit{\mu}°$ and $\mathit{\mu(a)}$ , i.e., Eq. $\mathit{\mu}3$.

Last but not least, adjactent phases to the phase under consideration also affect its surface potential, i.e., in the context of surface potentials, an additional standard condition must be defined. Usually, this tends to be the phase-vacuum interface (air, where surface potentials have been measured to date, is generally considered a good approximation of a vacuum).

However, it is not possible to measure $\mathit{\mu}$ without using extra-thermodynamic assumptions, which is why the IUPAC prefers $\mathit{\alpha}$ (at least for the $\mathrm{SHE_{H_{2}O}}$ , see above). On the other hand, it has been pointed out that “the assumption of additivity of dispersion and electrostatic forces has no thermodynamic basis”,^[6] or, in other words, the summation of $\mathit{\mu}$ and $\mathit{\varphi}$ (or $\mathit{\mu}$ and $\mathit{\chi}$ ) is also an extra-thermodynamic assumption. Thus, we currently see no convincing advantages for defining pH and pe scales with $\mathit{\alpha}$. However, for consistency reasons, both scales should be based on the same quantity, not only if one wants to continue electrochemical pH measurement. Therefore, the unifying $\mathrm{pH}_{\mathrm{abs}}^{\mathrm{H_{2}O}}$ and $E_{\mathrm{abs}}^{\mathrm{H_{2}O}}$ scales are based on $\mathit{\mu}$.

Literature

[1] G. Job, F. Herrmann, Eur. J. Phys. 2006, 27, 353.

[2] E. A. Guggenheim, J. Phys. Chem. 1929, 33, 842.

[3] M. Sastre, J. A. Santabella, J. Chem. Educ. 1989, 66, 403.

[4] https://goldbook.iupac.org/terms/view/S05917 (last assessed on 2026.03.04.)

[5] Y. Marcus, Ion Solvation, John Wiley, Chichester, 1985.

[6] B. A. Pethica, Phys. Chem. Chem. Phys. 2007, 9, 6253.

Publications

• “Proton solvation in water, methanol, ethanol, acetonitrile and methyl formate: A critical assessment of experimental and theoretical approaches to extend the solvent-independent unified acidity scale.” Regina Stroh, Niklas Gebel, Timo Kienzle, Valentin Radtke, and Ingo Krossing*, ChemPhysChem 2026, 27, e202500349, https://doi.org/10.1002/cphc.202500349
• “About a Trillion Times more Acidic than Expected? On the Determination of the Unified Acidity of Sulfuric Acid.” Valentin Radtke, Monika Bäuerle, Regina Stroh, Timo Kienzle, Daniel Himmel, Agnes Heering, Jaak Nerut, Enn Lust, Ivo Leito, Ingo Krossing*, Angew. Chem., Int. Ed. Engl. 2025, 64, e06429. https://doi.org/10.1002/anie.202506429
• “The Unified Redox Scale for All Solvents: Consistency and Gibbs Transfer Energies of Electrolytes from their Constituent Single Ions.” Valentin Radtke*, Denis Priester, Agnes Heering, Carina Müller, Thorsten Koslowski, Ivo Leito, Ingo Krossing*, Chem. Eur. J. 2023, 29, e202300609. http://dx.doi.org/10.1002/chem.202300609
• “Measurements and Utilization of Consistent Gibbs Energies of Transfer of Single Ions: Towards a Unified Redox Potential Scale for All Solvents.” Valentin Radtke*, Niklas Gebel, Denis Priester, Andreas Ermantraut, Monika Bäuerle, Daniel Himmel, Thorsten Koslowski, Ivo Leito, Ingo Krossing*, Chem. Eur. J. 2022, 28, e202200509. https://doi.org/10.1002/chem.202200509
• “A unified pH scale for all solvents. Part I – Intention and Reasoning.” V. Radtke*; D. Stoica; I. Leito; F. Camões; I. Krossing; B. Anes, M. Roziková; L. Deleebeeck, S. Veltzé; T. Näykki; F. Bastkowski, A. Heering; Nagy Dániel; R. Quendera; L. Liv, E. Uysal; N. Lawrence, Pure App. Chem. 2021, 93, 1049–1060. https://doi.org/10.1515/pac-2019-0504
• “The Inverted Philosopher´s Stone: How to turn Silver to a Base Metal.” V. Radtke*, K. Pütz, D. Himmel, I. Krossing, J. Solid State Electrochem. 2020, 24, 2847. https://doi.org/10.1007/s10008-020-04633-y
• “Symmetric potentiometric cells for the measurement of unified pH values.” Agnes Heering *, Daniela Stoica , Filomena Camões , Bárbara Anes , Dániel Nagy , Zsófia Nagyné Szilágyi , Raquel Quendera , Luís Ribeiro , Frank Bastkowski , Rasmus Born , Jaan Saame , Jaak Nerut , Silvie Lainela , Lokman Liv , Emrah Uysal , Matilda Roziková , Martina Vičarová , Alan Snedden , Lisa Deleebeeck , Valentin Radtke , Ingo Krossing , Ivo Leito *, Symmetry 2020, 12, 1150. https://doi.org/10.3390/sym12071150
• “The Experimental Determination of the Gibbs Energy of Transfer of Single Ions without Extra-thermodynamic Assumptions, Part II: Evaluation of the Role of Ion Solvation and Ion Mobilities on the Liquid Junction Potentials at an “Ideal” Ionic Liquid Salt Bridge. The System Ag⁺ in Water and Acetonitrile.” A. Ermantraut, V. Radtke, N. Gebel, D. Himmel, T. Koslowski, I. Leito, I. Krossing*, Angew. Chem., Int. Ed. 2018, 57, 2348–2352. http://dx.doi.org/10.1002/anie.201707334
• “The Experimental Determination of the Gibbs Energy of Transfer of Single Ions without Extra-thermodynamic Assumptions, Part I: The Operating Procedure for the Use of an Ideal Ionic Liquid Salt Bridge between Half Cells with Different Solvents.” V. Radtke, A. Ermantraut, D. Himmel, T. Koslowski, I. Leito, I. Krossing*, Angew. Chem., Int. Ed. 2018, 57, 2344–2347. http://dx.doi.org/10.1002/anie.201707333
• “Basic Remarks on Acidity.” (Review) Daniel Himmel, Valentin Radtke, Burkhard Butschke, Ingo Krossing*, Angew. Chem., Int. Ed. 2018, 57, 4386-4411. http://dx.doi.org/10.1002/anie.201709057
• “A Unified View to Brønsted Acidity Scales: Do We Need Solvated Protons?” E. Paenurk, K. Kaupmees, D. Himmel, A. Kütt, I. Kaljurand, I.A. Koppel, I. Krossing, I. Leito, Chem. Sci. 2017, 8 (10), 6964-6973. https://doi.org/10.1021/acsomega.6c02089
• “Absolute Brønsted Acidities and pH Scales in ILs.”  Daniel Himmel, Sascha Karl Goll, Franziska Scholz, Valentin Radtke, Ivo Leito, Ingo Krossing*, ChemPhysChem 2015, 16, 1428-1439. http://dx.doi.org/10.1002/cphc.201402906
• “The protoelectric potential map (PPM): an absolute two-dimensional chemical potential scale for a global understanding of chemistry.” V. Radtke, D. Himmel, K. Pütz, S. K. Goll, I. Krossing, Chem. Eur. J. 2014,20, 4194–4211. http://dx.doi.org/10.1002/chem.201302473
• “Bulk Gas Phase Acidity.” Daniel Himmel, Sascha K. Goll, Ivo Leito, and Ingo Krossing*, Chem. Eur. J. 2012, 18, 9333-9340. https://doi.org/10.1002/chem.201104025
• “Anchor points for the Unified Brønsted Acidity Scale: The rCCC Model for the Calculation of Standard Gibbs Energies of Proton Solvation in eleven representative Liquid Media.” Daniel Himmel, Sascha K. Goll, Ivo Leito and Ingo Krossing*, Chem. Eur. J. 2011, 17, 5808-5826. https://doi.org/10.1002/chem.201003164
• “A unified pH scale for all phases.” Daniel Himmel, Sascha K. Goll, Ivo Leito and Ingo Krossing*, Angew. Chem., Int. Ed. Engl. 2010, 49, 6885-6888. https://doi.org/10.1002/anie.201000252

Autoprotolysis of selected Solvents

All values are taken from Izutsu.^[1]

Literature

[1] K. Izutsu, Electrochemistry in Nonaqueous Solutions, Wiley-VCH, Weinheim, 2002.

Abbreviations of Solvents

Glossary