Article pubs.acs.org/jchemeduc
Electrolyte Solutions and Specific Ion Effects on Interfaces Ran Friedman* Department of Chemistry and Biomedical Sciences, Linnæus University, 391 82 Kalmar, Sweden, and Linnæus University Centre of Excellence for Biomaterials Chemistry, 391 82 Kalmar, Sweden ABSTRACT: Introductory general and physical chemistry courses often deal with colligative properties of solutions and do not discuss nonideal solutions in detail. Yet, a growing body of evidence reveals that even at physiological concentrations electrolyte solutions cannot be treated as ideal when a charged or partially charged solute (such as a protein) is present in the solution. In such cases, the interactions between the salt ions and the solute depend on the specific ions that constitute the electrolyte solution (specific ion effects). For example, the catalytic efficiency of an enzyme may be different in NaCl and KCl solutions. In this article, specific ion effects are reviewed from a historical perspective, and then the current state of knowledge is presented at a qualitative level that is appropriate for second-year or advanced undergraduate science students. Finally, the related nomenclature (Bjerrum ion pairs, Hofmeister series, lyotropic series, and specific ion effects) is analyzed, and some suggestions are made with respect to the terminology, to make it more accessible to students. The material is appropriate for courses where solution chemistry is discussed, for example, in physical chemistry. In addition, it may be included in the chemistry curriculum for life or pharmaceutical sciences. KEYWORDS: General Public, Upper-Division Undergraduate, Second-Year Undergraduate, Physical Chemistry, History/Philosophy, Textbooks/Reference Books, Biophysical Chemistry, Water/Water Chemistry
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Avogadro.8 Based on the same analogy, van’t Hoff coined the term ideal solutions for solutions that are diluted sufficiently that the volume of the solute is negligible with respect to the volume of the solvent. When discussing mixtures of solvents, solutions that obey Raoult’s law for any composition of the components (not only when one liquid is dilute) are referred to as ideal solutions in textbooks.6 A more general approach was devised by Edward Wight Washburn in 1910. Washburn suggested that colligative properties of solutions are governed by purely thermodynamic relations between the osmotic pressure, vapor pressure, gas and liquid molar volumes, and the temperatures of phase transition.9 Washburn referred to ideal solutions as pure mixtures in which no association or dissociation takes place, regardless of the concentration or the number of solutes. The term regular solutions was introduced by Hildebrand and refers to solutions for which the entropy of mixing is zero (as in ideal solutions), that is, the molecules of the solute and the solvent are distributed randomly, but the enthalpy of mixing is different than zero.10 Cantelo went one step further by introducing the chemical potential of a solute as a function of its activity.11 This is the level that is discussed today in introductory physical chemistry textbooks. Further developments, such as those dealing with the statistical mechanical theory of solutions, have emerged since the 1950s, but are generally of limited interest in undergraduate chemistry because of their mathematical complexity, although many studies involving such methods have been useful to infer on interesting chemical and physical observations.12−15
he fundamental work of Franz Hofmeister on the effect of salts on the precipitation of proteins1 is of great importance in biomolecular2 and colloid chemistry.3 Students of chemistry and life sciences normally encounter topics related to solutions in their first or second course of physical chemistry. Yet, a brief examination of several popular introductory physical chemistry textbooks4−6 revealed that the subject of the salt effects and precipitation is not covered in these books. The lack of coverage of the topic motivated me to write this article. Following the 29th International Conference on Solution Chemistry in 2005, the International Union of Pure and Applied Chemistry (IUPAC) solicited an article on specific ion effects from the German chemist Werner Kunz. After mentioning the Debye−Hückel theory, Kunz writes:7 The problems begin with the second-order ef fects, and, to be honest, the progress since the 1920s is not very impressive Although we are far from a full understanding of specific ion effects, which are second-order to pure charge effects, quite a few studies have been carried out since the 1920s (some of which are briefly reviewed by Kunz7). Chemistry graduates deserve to be aware of those effects that play an important role in solution chemistry. The material presented here could be introduced to students in physical chemistry courses that discuss solutions, prior to addressing nonideal solutions or following the discussion on colligative properties.
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GENERAL THEORIES OF SOLUTIONS
The foundations of solution theory are attributed to Jacobus Henricus van’t Hoff, who, in 1887, devised a theory of solutions analogous to the gas laws of Boyle, Gay−Lussac, and © XXXX American Chemical Society and Division of Chemical Education, Inc.
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ELECTROLYTE SOLUTIONS The discovery of Arrhenius that salts break into conducting species in water16 led to a distinction between electrolytic and nonelectrolytic solutions and to separate theories that dealt with each. It was soon realized that colligative properties of electrolyte solutions deviated from an ideal behavior at millimolar concentrations. Samuel Milner developed a theoretical framework to explain why the lowering of the freezing point of electrolytic solutions was less than expected from their colligative properties. Milner explained that ions in solutions orient themselves according to their charges; oppositely charged ions have a higher probability of being closer to each other, which generates a solution structure that is more organized.17 This description can also be understood in terms of electrolytic solutions having lower entropies compared to solutions containing the same number of uncharged particles of similar sizes. Following Milner’s ideas, a successful model to account for the deviation from expected (ideal) behavior of electrolytic solutions was developed by Debye and Hückel in 1923.18 The Debye−Hückel theory is still in use today for qualitative estimations of activity coefficients and electrostatic interactions. It is therefore covered in basic physical chemistry courses and textbooks. The extended Debye−Hückel equation reveals that the activity coefficients of the solution depend not only on the concentration of the ions but also on their charges and sizes. Niels Bjerrum modified the Debye−Hückel theory to account for ions that are not fully dissociated in solution,19 as is the case for virtually all complexes of multivalent ions. Bjerrum described all ions within a certain cutoff distance rc as associated, where the cutoff depends on the charge of the ions and the dielectric permittivity of the solutions. Ions that are located further apart than rc are considered unassociated and contribute to the bulk properties of the solution. Bjerrum’s treatment, though less well-known, is of great significance in marine and colloid chemistry. Although accounted for only in specialized textbooks,20,21 it presents an elegant approach to one of the limitations of the Debye−Hückel treatment and should perhaps be considered for a discussion in undergraduate courses.
chemistry (the composition of marine aerosols), and inorganic chemistry (interfaces of ionic liquids, interactions between solutions and inorganic surfaces). A common explanation for the Hofmeister series, which is still found in some textbooks, is that specific ion effects are due to interactions with the water. Kosmotropic (order-making) ions such as SO2− are strongly hydrated. Such ions, therefore, organize the water located within 1−2 solvation shells around them and decrease the solubility of the proteins, which in turn drives their precipitation. Chaotropic (order-breaking) ions, such as SCN−, have the opposite effect. They interact weakly with water molecules, making the nearby water less ordered. The disordered water molecules can more readily interact with proteins and other large solutes, which increases their solubility. The ordering of water molecules, however, cannot explain some of the observed phenomena. For example, whereas chaotropic and kosmotropic effects can explain the influence of the electrolyte on the molar activity coefficient of protons and many inorganic cations and anions,23,24 this explanation in not valid for acetate ions in water,23 which are often considered as models for carboxylates in proteins. Moreover, sensitive spectroscopy measurements do not reveal any long-range high-order effects of water.25 An alternative explanation relates the Hofmeister series to specific interactions with the surface, for example, between Na+ or K+ ions and carboxylates.26 Such effects are manifested to a greater degree at lower ionic strengths,27 where screening of electrostatic interactions by the bulk solution is less pronounced. A useful concept in this respect, adapted here from Marcus,28 is that of the average distance between ions. Obviously, the more concentrated the solution the more likely it is to find a counterion near any given ion. In the case of a macromolecule or surface that has many charges attached to it (e.g., a polyelectrolyte gel), the higher the concentration of the solution, the more likely it is that two charges (a charged surface group of the macromolecule or polyelectrolyte and an ion) interact purely by chance. It is very probable that ions might be found within 1 nm of the protein surface residues in a 1 M solution (but this is not the case at physiological concentrations). Accordingly if a charged chemical group interacts with an oppositely charged ion in a 1 M solution, it is likely to interact also with others.29 However, care should be taken not to over interpret this structural concept. According to Coulomb’s law, the electrostatic interaction energy between two charged particles, E, is
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SPECIFIC ION EFFECTS: THE HOFMEISTER SERIES A few years after the discovery of Arrhenhius, Franz Hofmeister and his colleagues reported that various salts had a different tendency to precipitate proteins is solution.1,22 For example, it takes a small amount of sodium sulfate to precipitate lysozyme in water, but a much higher amount of sodium chloride, and an even higher amount of sodium nitrate is necessary to observe precipitation (note that Hofmeister reported on the mass quantity of salt. i.e., grams rather than moles, but the order prevails if one uses molar amounts). Interestingly, Hofmeister’s findings could be repeated with different proteins; salts that had the tendency to precipitate or (conversely) solubilize one protein had the tendency to precipitate or solubilize other proteins as well. These effects appeared only at high concentrations (C > 1 mol dm−3 = 1 M), and their origin is still a mystery. Hofmeister effects (also known as lyotropic effects, see below) are widely recognized in surface chemistry today, although they are seldom mentioned in introductory physical chemistry textbooks. Specific ion effects are not limited to proteins, and may involve other solutes in electrolyte solutions. They thus play a role in biochemistry (enzymology, macromolecular interactions), marine chemistry, atmospheric
E=
q1q2 rεr
(1)
where q1 and q2 are the ionic charges and εr is the dielectric coefficient of the solvent. The use of SI units is inconvenient in this case, as this implies that the distances are given in meters. It is therefore more common to use a conversion factor of 33.230 and express the distance r in nm and energy in kCal/mol: E = 33.2
q1q2 rεr
kCal/mol
(2)
For a 1 M, 1:1 solution, the number of molecules ñ per nm3 is n ̃ = 6.02
B
1023L−1 1021
nm 3 L
(3)
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The mean distance between any two ions is (1/ñ)1/3, that is, 1.18 nm. Thus, according to eq 2, |E| = 0.36 kcal/mol, or about 0.6 kT at room temperature. Yet 1 M solutions are not ideal dilute solutions, in spite of the weakness of ion−ion interactions. Nonideality is explained by interactions with the solvent (note also that a large fraction of the oppositely charged ions are closer than the average distance at any given time point).
softness have all been used to explain the preference of certain ion pairs based on the principle of matching properties. Ion Binding to Surfaces May Depend on Local Geometry
Unlike the surface of water, the surface of proteins, and many colloidal solutes, is irregular and somewhat rugged; it includes protrusions and crevices. The local geometry around a residue as well as its immediate neighbors can make the residue more (or less) prone than the average to bind oppositely charged ions.42 Consider, for example, a small crevice on the protein surface, not much larger than an hydrated ion. Ions will be less likely to form contacts with residues located inside such crevices, but also less likely to leave, simply due to the geometry. Moreover, the local electrostatics will also play a role. Ions will be more likely to be located in areas on the surface where the overall electrostatic potential is opposite in sign. Local geometry is of smaller significance at molar ionic strengths, because the protein surface is saturated with ions, but can be manifested at physiological conditions.
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LOCAL ION EFFECTS ON SURFACES Even if the local effects of ions on surfaces cannot explain the Hofmeister series, they are of great importance, and some qualitative understanding of the underlying principles can be of benefit to chemistry students. Local ion effects are evident when a charged surface is present in an electrolyte solution. The higher the ionic strength of the solution, the weaker are the local ion effects because the electrostatic field around the solute is reduced,31 making local ion effects evident in ambient, physiological, and standard (molar) concentration solutions,27 but not at much higher concentrations. Local ion-surface effects can be thought of in terms of pairing of solute and salt charges, such as charged protein residues and oppositely charged ions, or DNA and positive salt ions. The preference to specific ions can be explained by matching properties that would make two charged entities more likely to form an ion pair. Moreover, in the case of asymmetric solutes such as proteins or polyelectrolyte gels, the local geometry of the solute surface will render some regions more attractive to ions than others, as explained below.
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TERMINOLOGY At this point, a discussion on the terminology of specific ion effects may be warranted. Many students find chemistry complex.43−45 Several factors contribute to this, and the use of specialized terminology is certainly one of them. Indeed, scientific terminology is a stumbling block for students of many disciplines, including chemistry,46,47 physics,48 life sciences,49 and medical sciences.50 The use of multiple terms with similar meanings or multiple meanings for a single term is detrimental in this respect. The need for standardization in chemistry was recognized more than a century ago and contributed to the formation of IUPAC. One of IUPAC’s publications, the Gold Book (http://goldbook.iupac.org/), is a compendium of chemical terminology. The Gold Book includes the following definition for an ideal dilute solution: Dilute solution in which the solute may be regarded as obeying Henry’s law, so that all the solute activity coef f icients may be approximated to 1. To reduce misconceptions among solution chemistry students, it is perhaps better to refrain from using the term ideal solution in class and refer to ideal dilute solutions only. Although the term Bjerrum ion pair does not have an entry in the Gold Book, it is discussed under ion pairs: Following Bjerrum, oppositely charged ions with their centers closer together than a distance:
Ion Pairing
The association of ions can influence the properties of the solution, as identified by Bjerrum. However, Bjerrum’s model did not account for molecular models of association, and his theory is sometimes criticized due to this simplification. In contemporary solution chemistry,32,33 there is a distinction between contact ion pairs (CIP) and solvent-separated ion pairs (SIP). To be considered as an ion pair, two ions of opposite signs must be within a given distance for a time period longer than that needed for diffusion over the same distance.34 Ion pairing plays an important role in ion-surface interactions and affects the behavior of polyelectrolyte solutions (e.g., polysaccharide gels), proteins, and membranes. Moreover, specific ion pairing can determine how the solvation of ionic liquids in water is affected by salt.35 Protein stabilization can also be rationalized by ion pairing, whether it occurs through interactions with salt ions36 or between oppositely charged amino acid residues.37
rc =
Matching Properties
8.36 × 103z+z − εr T
are considered to constitute an ion pair (’Bjerrum ion pair’). [z+ and z− are the charge numbers of the ions, and εr is the relative permittivity (or dielectric constant) of the medium.] (note that the symbol rc is used for the distance in nm, to be consistent with the equations above). The terms Hofmeister effect, Hofmeister series, lyotropic effect, lyotropic series, and specific ion effects are not defined in the Gold Book and warrant further discussion. Hofmeister himself was not familiar with the concept of ions when he started working on protein precipitation by salt, and hence terms such as the “Hofmeister series of anions” may be somewhat misleading. Moreover, deviations between the ordering of ions according to the Hofmeister effect and their ability to bind to proteins were noted by Loeb already in 1920,51 and are now used for the development of anion-specific
The pairing between two charged particles may be governed by physical properties that match each other. As an example, two oppositely charged ions that have similar hydration free energies will have a higher tendency to partially lose their solvation shells and form ion pairs. This is the reason for the higher ability of Na+ ions to bind to carboxylates at the protein surface compared with K+.38 Likewise, oppositely charged ions that have the same affinity to the air/water interface can associate, even when they are different in sizea phenomenon that can be detected at micromolar concentrations.39 Pearson’s Hard−Soft Acid−Base (HSAB) theory40 has also been used to explain ion pairing in solution,41 suggesting that ions with similar hardness will be more likely to pair. Thus, hydration free energies, affinities to the ion−water interface, and hardness or C
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sensors or separators.52,53 The adjective “lyotropic” is used by IUPAC only with respect to a mesophase. Oxford’s English Dictionary54 refers both to the lyotropic series and to the lyotropic mesophase; the lyotropic series is defined as a series where ions are arranged according to their tendency to precipitate a lyophilic (=dispersed colloidal) solute. Since the term lyotropic is used in another context, it is better to avoid confusion by using a different term. Specific ion effects is perhaps the most general description for interactions involving ions that go beyond their charges,55 or, as recently stipulated by Lo Nostro and Ninham:56 ...by specif ic ion ef fects we mean ef fects not accommodated by classical theories of electrolytes. This is also the simplest term, as it does not involve the concepts of lyotropy or precipitation, and therefore it may be preferred over the Hofmeister/lyotropic effect.
relates student learning and development to exposure to the right subject matter. Individual courses are evaluated according to their content, and the proper content is therefore of utmost importance. The resource theory, on the other hand, relates student learning and development to the available resources: physical facilities, high quality personnel, and fiscal means. From this point of view, the professional educators should be well-trained, and it is important that they have sufficient knowledge. Because a teacher cannot be expert in everything, topics of importance should be covered by textbooks. The literature and teachers are both resources that should enable the learning. The newer theory by Astin, the theory of student involvement, puts the activity of the student in focus. Astin writes:58 On a more subtle level, the theory of student involvement encourages educators to focus less on what they do and more on what the student does: how motivated the student is and how much time and energy the student devotes to the learning process. The theory assumes that student learning and development will not be impressive if educators focus most of their attention on course content, teaching techniques, laboratories, books, and other resources. With this approach, student involvement - rather than the resources or techniques typically used by educators - becomes the focus of concern. The theory sees the actual content of the course as less important than the involvement of the teacher in motivating the students. Nevertheless, in my opinion it is still necessary that the study material will be up-to-date, interesting, and clear. Without this, one cannot expect the students to be motivated. The binding of the content to current research and development is, therefore, important, and the right content should be at least accessible to the students, if not taught explicitly. It is also worthwhile to use coherent terminology that can improve both understanding of the student and their ability to read the scientific literature.
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SOME COMMENTS ON COURSE CURRICULUM IN CHEMISTRY An interesting report related to curriculum development in natural sciences came from the American Association for Advancement of Science (AAAS) in 2000. Entitled “The Liberal Art of Science: Agenda for Action”, the AAAS study group discussed the level of scientific understanding that was required for undergraduate education in the 21st century. The study group argued that education in science should provide the students with a knowledge base that would enable them to educate themselves about contemporary science and technology, and with an understanding of the methods and processes of scientific inquiry. With this in mind, it is evident that education in chemistry should go beyond 19th and early 20th century models and be closely linked to today’s research and technology. I made two recommendations that stem from these: discussion of Bjerrum’s ions pairs in solution chemistry and teaching of specific ion effects and gave examples from contemporary research to justify the inclusion of the latter. Claxton57 describes three modes for the development of course curriculum: content-based, knowledge-based, and development model. The first two of these models are relevant to natural sciences, whereas the latter is more appropriate for applied subjects such as nursing or engineering. It is clear that proper content is crucial in a content-based model and Claxton writes: Such an approach can be appropriate if the course is responsive to the needs of the potential learners, if the faculty members teach so that the content is relevant to learners. Content-based curriculum development is, however, criticized when it leads simply to telling facts and not enabling learning. The proper approach, according to Claxton, is to ask “What can I do to help learners develop the abilities they need?”. Getting back to the AAAS report above, one can reply that science should be thought in a liberal education frameworkmentioning the facts but enabling the students to question the methods and conclusions. A knowledge-based approach emphasizes the knowledge from the perspective of the learner. Here, the contents depends on the objectives of the course. According to the AAAS report, science education should promote scholars that are able to deal with the challenges of the 21st century and for this it is necessary to include relevant content that is manifested in scientific research and development. Course development can also be discussed from the point of view of traditional and modern pedagogical theories.58 The subject-matter theory puts a great emphasis on the content. It
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CONCLUSIONS
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AUTHOR INFORMATION
This article gives a brief historical perspective to theories dealing with electrolyte solutions and discusses specific ion effects. Specific ion effects are the main focus, owing to their importance and lack of coverage by introductory textbooks. Some suggestions are made with respect to teaching of the subject. It is postulated that Bjerrum’s theory of associated ions, specific ion effects, and their causes should be discussed, or at least presented, as extra reading material in courses or textbooks that cover basic solution chemistry. Few remarks are made with respect to the terminology: (1) the term “ideal dilute solutions” should be used (following IUPAC), whereas the term “ideal solutions” should be avoided, (2) The term “Bjerrum ion pair” may be introduced, and (3) the term “specific ion effects” should be used for interactions that are not explained by the classical theories of electrolytes, such as the Debye−Hückel theory.
Corresponding Author
*E-mail:
[email protected]. Notes
The author declares no competing financial interest. D
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ACKNOWLEDGMENTS
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The author wishes to thank Karina Adbo for critically reading the manuscript, and Henrik Hegender, Elisabeth Elmeroth, two anonymous reviewers, and the scientific editor for comments on the manuscript. Funding from the Faculty of Sciences and Engineering at Linnæus University (project title “Metal Ions in Life”), the Carl Tryggers Foundation (project no. CTS 11:146), and the Swedish Royal Academy of Sciences (project no. FOA12V-021) for related research projects is also acknowledged.
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Journal of Chemical Education
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dx.doi.org/10.1021/ed4000525 | J. Chem. Educ. XXXX, XXX, XXX−XXX
