In the Classroom edited by
Concepts in Biochemistry
William M. Scovell Bowling Green State University Bowling Green, OH 43403
Molecular Handshake: Recognition through Weak Noncovalent Interactions
W
Parvathi S. Murthy Department of Chemistry and Biochemistry, Georgian Court University, Lakewood, NJ 08701;
[email protected] After nearly a century of anatomical study of geckos and their ability to adhere to and move rapidly along vertical walls and ceilings, the mechanism of this dry adhesion has been recently determined by Autumn et al. (1). The Tokay gecko’s foot has nearly five thousand keratinous hairs, each of which contains hundreds of projections terminating in a spatulashaped structure. The size and shape of these millions of hydrophobic tips or setae lead to remarkably strong adhesion to any surface (hydrophilic or hydrophobic) by van der Waals interactions. Collectively, this adhesion is strong enough to hold the gecko in place on a wall or ceiling. However, the interaction of each seta with the surface is weak and can be broken easily and gradually, nearly one interaction at a time, allowing the animal to move on. This example illustrates the significance of noncovalent interactions—ability to be formed and broken easily without much expenditure of energy—affecting the behavior of living and nonliving matter. The noncovalent interaction between molecules—the handshake in the form of electrostatic interactions: van der Waals interactions, hydrogen bonding, or charge transfer interaction—can be found in all matter. Among these, the weak interactions (some electrostatic interactions, van der Waals interactions, and hydrogen bonding) are of particular importance because of the ease with which they can be formed and broken. These interactions among molecules of a homogeneous substance are responsible for determining their bulk properties: melting point, boiling point, viscosity, surface tension, and so forth. The interactions between molecules of two different substances are significant in influencing the rates of chemical reactions, effectiveness of chromatographic separations, and molecular recognition in biological processes. The knowledge of noncovalent interactions between substances is utilized in the development of new materials and commercial products with specific utility. This article traces the development of our thinking about weak noncovalent interactions, highlights their salient features, and suggests ways for a comprehensive exposition of the principles of these noncovalent interactions in undergraduate chemistry curricula. History of Noncovalent Interactions Our present ideas of interaction between chemical substances have emerged over centuries. In 1743, Clairault envisioned interparticle forces between water and glass and between particles of water to explain capillary rise of water in a narrow glass tube (2). The idea that the properties of matter were related to interaction between atoms or molecules started emerging in the 19th century with the derivation of 1010
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the ideal gas equation of state
P V = N kB T
(1)
where P is pressure, V is volume, N is the number of molecules, kB is Boltzman’s constant, and T is the absolute temperature. This derivation assumed that gas molecules had negligible volume (point molecules or atoms) with no interaction between them. Although this implied that solids, liquids, and nonideal gases could possess interactions between atoms or molecules, it was late 19th century when the fact that molecules repel at short ranges and attract at long ranges was established (2). In 1873, J. D. van der Waals, after studying a number of real gases that exhibited deviation from the ideal gas equation under certain conditions, successfully modified the eq 1 to P +
a V
2
V − b
= N0 kBT
(2)
– where V is molar volume and N0 is Avogadro’s number (3). The parameter a relates to the strength of intermolecular attractive interactions and b relates to the effective volume of the spheres (repulsive interactions). Van der Waals considered atoms or molecules of real gases to be hard spheres that had certain extensions in space (not mass points) and even as neutral species attracted each other at long range. Equation 2 reduces to the ideal gas equation when a and b are zero. There was no clear understanding about the nature of these universal molecular attractions, but eq 2 provided for the first time quantitative information about the interactions. Early in the 20th century, Reinganum (1903), Debye (1912), and London (1927) contributed to a better understanding of intermolecular forces (2). Reinganum made an attempt to relate the interactions to the structure of molecules by proposing the presence of a “bipole” (a pair of opposite charges separated by a constant distances rigidly located within a molecule). This was strengthened by early investigations by Debye on permanent dipoles. Soon it became apparent that not all molecules had permanent dipoles. London’s publication on covalent bonding in H2 molecule led to an understanding of the chemical bonding within this molecule and the repulsive forces that operate at short distances. Following this, London also provided the first theory of the long-range attractive interactions (known as London forces) using quantum mechanics. The presence of a weak bond between certain molecules, involving hydrogen (hydrogen bonding), was proposed in 1919 and 1920 (4).
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Nature of Noncovalent Interactions In all matter, when atoms, ions, or molecules (of the same or different substances) approach each other to a proximity of ∼1 nm, there will be negligible overlap of electron clouds.1 However, there can be noncovalent attractive interactions between them. The various types of noncovalent interactions are: (i) ion–ion, (ii) ion–dipole, (iii) ion–induced dipole, (iv) dipole–dipole (v) dipole–induced dipole (vi) London dispersion forces (induced dipole–induced dipole interactions), (vii) hydrogen bonding interactions, and (viii) charge transfer complexes. The noncovalent interactions among substances with no net charge, iv–vi, are often collectively termed van der Waals forces. The nature and extent of noncovalent interactions (i–viii) is determined by the type of interacting species. The interaction energy2 depends on the nature of species involved, their relative orientation, and their separation. In the case of solids (macroscopic or microscopic objects or molecular species), surface uniformity (on a nanometer scale) will determine the extent of “contact” and hence the extent of noncovalent interactions between them (1, 5). In liquid, solution, or gaseous phase, there is a better chance of approach to such close distances and hence better interaction. The above interactions can also be characterized as weak or strong noncovalent interactions. Weak interactions are those that can be disrupted by thermal kinetic energy of the system3 and are transient (lifetime 10᎑12 s) in nature (6). The strength of these interactions under the existing conditions, taken collectively over the bulk of the substance, influences the behavior of this form of matter. Strong interactions (ion– ion and charge transfer types4) require changes in conditions such as chemical reagents, pH, temperature, or ionic strength
to be disrupted. This article is limited to the discussion of weak noncovalent interactions: van der Waals interactions and hydrogen bonding. The nature of weak noncovalent interactions is described in Table 1. When chemical species approach each other closely, there will be London forces of attraction between them. In addition, there may also be other types of weak interactions depending on the nature of the interacting species. Although, individually weak and operating over brief periods of time, collectively this form of molecular recognition mechanism can significantly influence the properties of chemical and biological systems. Through differential interaction with the reactants, products, and the transition state, solvents influence the rate of chemical reactions. Solvent interactions can also alter the mechanism and hence the products of the reaction. In chromatography, the stationary phase and the mobile phase compete to interact with the analytes effecting their separation. While the nonspecific but collective effect was significant in the last two examples, specific noncovalent interactions are responsible for enzyme catalysis, allosteric regulation of gene expression, and signal transduction processes in biological systems. The ease of formation and cleavage of intermolecular forces between a protein (enzyme) and its substrate (catalysis), a protein and its effector ligand (allosteric regulators), or a receptor protein and its specific signal molecule (signal transduction pathway) leads to establishment of a dynamic equilibrium between the “free” and the “bound” states of the protein. The existence of this equilibrium is crucial to the biological function of the protein. The availability of the binding substance (substrates, effectors, or signal molecules) shifts the equilibrium between the free (inactive) and bound (active) states of the protein and thus determines the its biological activity under the ex-
Table 1. Nature of Weak Noncovalent Interactions Types of Interacting Substances, A and B
Nature of Interaction: (a) Type of Force (b) Interaction Energya
ion–dipole
A: Ionic B: Polar covalent
ion–induced dipole
A: Ionic or polar covalent B: Nonpolar covalent
(a) Electrostatic (b) 0.2–4.0 kJ/mol (at separation distance of ~1 nm). Strength falls of with r ᎑2 for ion–dipole and r ᎑4 for ion–induced dipole interactions.
dipole–dipole
A, B: Polar covalent
dipole–induced dipole
A: Polar B: Nonpolar covalent
induced dipole–induced dipole (London forces, Dispersion forces)b
All substances
(a) Attractive interactions between fluctuating charges. (b) 0.2–2 kJ/mol (at separation distance of ~1 nm). Strength falls of with r ᎑6.
A: Compounds with strongly polar covalent bonds containing hydrogen atom as part of the dipole B: Compounds with lone pair electrons or negative charges or pi electrons
(a) Donor–acceptor/electrostatic interaction (b) 10–40 kJ/mol (at separation distance between A and B < 0.4 nm) (~2 kJ/mol for those involving pi electron acceptors).
Noncovalent Interaction
Electrostatic
van der Waals (substances with no net charge)
Hydrogen Bonding Requires interaction between specific atoms, as shown (A is the donor, B is the acceptor) A–H + :B → A–H:B Weakens A–H bond. Distance between H and B in the complex can be lower than sum of their van der Waals radii. Directional (usually 180⬚). Stoichiometric.
(a) Electrostatic (b) 0.5–4 kJ/mol (at separation distance of ~1 nm). Strength falls of with r ᎑3 for dipole–dipole and r ᎑6 for dipole–induced dipole interactions.
a The magnitude depends on the net charge, size of the species, separation distance (r), and temperature. Its value is inversely proportional to rn (n varying from 2–6 for various interactions). The values indicated here are typical values used in common text books. bMajor contributor to noncovalent interactions except in small, highly polar molecules. Often referred to as hydrophobic interactions, in biological systems.
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isting condition. The weak intramolecular interactions within the protein, nucleic acid, and carbohydrate chains support retention of the gross three-dimensional structures of the biopolymers while allowing for regional conformational flexibility necessary for the biological function. A biological cell membrane is a phenomenal structure that is held together only by noncovalent forces. This structure gives the membrane fluidity and determines its functions, such as self-assembly, self-healing, and the ability to separate contents while allowing specific substances to cross it. Noncovalent Interactions in the Undergraduate Chemistry Curriculum One of the main goals of chemical education is to teach students to understand the relationship between molecular structure or composition of matter and its properties. The relationship between structure and property (established by studying known matter or processes) is the tool that can lead to the design of new materials and processes. Much of the matter we come across in everyday life is not an individual atom or molecular species but is heterogeneous in nature. To gain an insight into the structure–property relationships requires a good understanding of the structure or composition of matter (chemical bonding) and the nature of noncovalent interactions between molecular species as they exist in close proximity to one another. Such an important universal concept deserves greater emphasis in chemistry curriculum. The following paragraphs describe the extent of coverage of noncovalent interactions in the current undergraduate chemistry curriculum. A significant portion of the general chemistry curriculum is allocated to teaching covalent or ionic bonding interactions that make up individual molecules and network structures. Although matter is usually not found as single atomic or molecular species, the textual resources allocated to the topic of interactions between molecules or their influence on properties of matter is limited to pure substances. Some general chemistry textbooks show the following coverage of noncovalent interactions (7): The kinetic theory of gases (in the chapter on gases) includes a statement such as attractive and repulsive interactions between gas molecules are negligible in the gaseous state. There is little, if any, description of what these forces are. The chapter also includes a section on van der Waals’s equation and PV兾RT versus P behavior of real and ideal gases. Usually, following this chapter is the chapter on intermolecular forces in pure solids and pure liquids with definitions of various noncovalent interactions and their influence on physical properties: surface tension, viscosity, and boiling point. The chapter on properties of solutions describes the influence of the nature of solutes and solvents on dissolution with a few typical examples. The chapter includes the effect of solutes on colligative properties of solvents in a quantitative manner without delving into the nature of solute–solvent interactions that lead to the colligative properties. The chapters on intermolecular forces and properties of solutions also, generally, fall into chapters 10–14 of the textbook—towards the end of the semester—and may not get full coverage. The chapter on reaction kinetics rarely mentions the importance of the role of solvents in influencing the kinetics or mechanisms of chemical reactions. The laboratory experi1012
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mental content may include an experiment on the influence of the nature of solutes and solvents on dissolution (8). Typical organic chemistry textbooks often introduce the topic on noncovalent interactions in early chapters and briefly describe their influence on solubilities of compounds with specific functional groups in water (9). The laboratory synthesis experiments do not incorporate the importance of the choice of solvents in these reactions. The importance of noncovalent interactions in determining the effectiveness of chromatographic separations is briefly included, if at all, even though this technique could make up a significant portion of the laboratory experiments. The chapters on carbohydrates, proteins, nucleic acids, and lipids refer to noncovalent interactions as they describe the structure of these compounds. The fact that the biological activity of these molecules, very often, is entirely a consequence of their three-dimensional structure, is not emphasized. Physical chemistry and analytical chemistry courses may include spectroscopic methods to study hydrogen-bonding interaction but is limited to few specific systems, for example, keto–enol tautomerization of acetylacetone (10). Biochemistry texts describe these interactions as contributing to structural aspects of biomolecules and do not describe the importance of the weak interactions in binding processes—ease of formation and breakage (11). There have been nearly fifty publications in this Journal in the past ten years on topics such as molecular recognition or intermolecular forces. Most of these articles describe experiments pertaining to one specific system, involving hydrogen-bonding interactions, or are mostly directed to upper-level (physical chemistry or biochemistry) laboratory instruction. Several recent articles in this Journal have identified the need for better instruction in this field; Rebek et al. have also identified the lack of sufficient time for teaching noncovalent interactions in undergraduate chemistry curricula (12). Henderleiter et al. have discussed how students in organic chemistry classes, while able to give definition and even identify hydrogen bonding in substances, had difficulty relating the influence of this interaction on the properties of substances (13). They point the cause to the confusion in the descriptions found in general chemistry and organic chemistry textbooks. While discussing the teaching of noncovalent interactions in a biochemistry course, Cox has identified the shortcomings of biochemistry textbooks in describing the various noncovalent interactions (14). In conclusion, the concepts of noncovalent interactions are included in a sporadic way in various chemistry classes but the bigger picture of the nature of molecular recognition through weak noncovalent interactions, universality (when any substance comes in contact with any other substance) and the significance of their weakness or transient nature, is not portrayed in the curriculum. Considering all matter is “chemical”, chemistry students should be able to identify the type of interactions between any two “objects” (macroscopic, microscopic, or molecular). They should be able to suspect noncovalent interactions when they see a paint sticking to one surface and not to another. To provide this broader perspective, in addition to emphasizing the content that is already in current textbooks, it is important to reinforce the concept of noncovalent interactions throughout the undergraduate curriculum.5 The main concepts, such as the nature of chemical reactions, dynamic equilibrium, methods of
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separation, and bulk properties, are an integral part of every chemistry course. Utilizing a lecture to bring together the concepts of noncovalent interactions encountered throughout the semester, in different contexts, would be time well spent. Examples such as the gecko’s walk illustrate these principles well. Laboratory instruction on experiments in which the students are using mixtures is a good opportunity to provoke thinking about the interactions (which substance is recognizing which other substance and to what degree) and any possible effects of these interactions. All chemistry students must be aware of the value of chemical principles in expanding knowledge in other disciplines. The influence of noncovalent interactions on biological processes and materials research and development are good examples of such applications. A description of some examples is available in the Supplemental Material .W Summary Molecular recognition through noncovalent interactions, as an interdisciplinary science, is contributing enormously to chemical genomics, drug discovery, and materials design (16, 17). New methods of detection and quantification of these interactions are developed at an increasing pace (18). Many students will go on to industrial or academic research laboratories to pursue knowledge in this area. Noncovalent interactions and the consequent molecular recognition should be an important component of undergraduate chemistry curriculum. Acknowledgment The author is grateful to Jannette L. Carey, Department of Chemistry, Princeton University, Princeton, NJ, for helpful discussions. WSupplemental
Material
Examples of the influence of weak noncovalent interactions on chemical and biological systems with references to literature are available in this issue of JCE Online. Notes 1. The van der Waals radii of nonmetallic atoms range from 0.12 nm for hydrogen to 0.22 nm for iodine. 2. Spectroscopic methods (IR and NMR), diffraction (X-ray and neutron diffraction), thermochemical (calorimetry of heats of mixing or dilution, determination of bulk properties, enthalpies or determination of equilibrium constants for the formation of nonbonded complexes), and computational methods are used for qualitative and quantitative study of these interactions. 3. For gaseous molecules, at 300 K, the thermal energy is of the order of ∼3.7 kJ兾mol (5). 4. Interaction energy of ion–ion interactions is of the order of ∼300 kJ兾mol and charge transfer interactions is of the order of ∼50 kJ兾mol. 5. Recent textbooks have adopted a different approach to introducing noncovalent interactions: introducing the concept in earlier chapters, including examples from biological chemistry along with those of small molecules or dedicating a chapter to molecular recognition (15).
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Literature Cited 1. Autumn, K.; Sitti, M.; Liang, Y. A. Proc. Natl. Acad. Sci. U.S.A. 2002, 99, 12252–12256. 2. Kaplan, I. Theory of Molecular Interaction. In Studies in Physical and Theoretical Chemistry; Elsevier: Netherlands, 1986; Vol. 42, pp 1–12. Margenau, H.; Kestner, N. R. Theory of Intermolecular Forces, 2nd ed.; International Series of Monographs in Natural Philosophy, Vol 18; Pergamon: Oxford, 1969; pp 1–14. 3. Chu, B. Molecular Forces; Interscience: New York, 1967; pp 1–15. 4. Jefferey, G. A. An Introduction to Hydrogen Bonding; Oxford University Press: New York, 1997; pp 1–3. 5. Rigby, B. M.; Smith, E. B.; Wakeham, W. A.; Maitland, G. C. The Forces Between Molecules; Clarendon Press: Oxford, 1986; pp 62–80. 6. Israelachvili, J. N.; Tabor, D. Proc. R. Soc. London, Ser. A 1972, 331, 19–38. 7. Hill, J. H.; Petrucci, R. H.; McCreary, T. W.; Perry, S. S. General Chemistry, 4th ed.; Prentice Hall: New York, 2005. McMurry, J. M.; Fay, R. C. Chemistry, 4th ed.; Pearson Prentice Hall: Upper Saddle River, NJ, 2004. Chang, R. Chemistry, 7th ed.; McGraw Hill: New York, 2004. Atkins, P.; Jones, L. Chemical Principles, 2nd ed.; W. H. Freeman: New York, 2002. Ebbing, D. D.; Gammon, S. D. General Chemistry, 8th ed.; Houghton Mifflin: New York, 2005. 8. Neidig, H. A.; Spencer, J. N. Solutions. In Modular Laboratory Program in Chemistry; Neidig, H. A., Ed.; Chemical Education Resources, Brooks/Cole: Montery, CA, 2003. 9. Wade, L. G., Jr. Organic Chemistry, 6th ed.; Pearson Prentice Hall: Upper Saddle River, NJ, 2006. Carey, F. A. Organic Chemistry, 6th ed.; McGraw Hill: New York, 2006. Vollhardt, K. P. C.; Schore, N. E. Organic Chemistry: Structure and Function, 4th ed.; W. H. Freeman: New York, 2003. 10. Sawyer, D. T.; Heineman, W. R.; Beebe, J. M. Chemistry Experiments for Instrumental Methods; Wiley: New York, 2003; pp 297–300. Frohlich, H. J. Chem. Educ. 1993, 70, A3–A6. 11. McKee, T.; McKee, J. R. Biochemistry: The Molecular Basis of Life, 3rd ed.; McGraw Hill: New York, 2003. 12. Hof, F.; Palmer, L. C.; Rebek, J., Jr. J. Chem. Educ. 2001, 78, 1519–1522. 13. Henderleiter, J.; Smart, R.; Anderson, J.; Elian, O. J. Chem. Educ. 2001, 78, 1126–1130. 14. Cox, J. R. J. Chem. Educ. 2000, 77, 1424–1428. 15. Chemistry: A Project of the American Chemical Society; Bell, J. A., Ed.; W. H. Freeman: New York, 2004. Silverberg, M. S. Chemistry: The Molecular Nature of Matter and Change, 4th ed.; McGraw Hill: New York, 2006. Fox, M. A.; Whitesell, J. K. Organic Chemistry, 4th ed.; Jones and Bartlett; Boston, 2005. 16. Handbook of Polyeletrolytes and Their Applications, Vol. 3; Tripathy, S. K., Kumar, J., Nalway, H. S., Eds.; American Scientific: Los Angeles, CA, 2002. 17. Protein Ligand Interactions: From Molecular recognition to Drug Design; Bohm, H.-J., Schneider, G., Eds.; Methods and Principles in Medicinal Chemistry, Vol 19; Wiley-VCH: Weinheim, Germany, 2003. Sharma, C. V. K. J. Chem. Educ. 2001, 78, 617–622. 18. Biological NMR Spectroscopy; Markley, J., Opella, S. J., Eds.; Oxford University Press: New York, 1997. Real-Time Analysis of Biomolecular Interactions: Applications of BIACORE; Nagata. K., Handa, H., Eds.; Springer-Verlag: New York, 2000.
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