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Hydrophobic Interactions

We have studied the role of the hydrophobic effect (involving the favorable entropic release of caged water molecules about solvent-exposed hydrophobic groups) in driving micelle and bilayer formation. Does this also drive protein folding? To explore this question, we will study the thermodynamics of small nonpolar molecules, especially benzene, with water and ask whether the thermodynamic parameters associated with benzene solubility are similar to those associated with protein stability. If this analogy holds, anything that will promote benzene solubility will lead to increased hydrophobic amino acid side chain exposure to water and hence protein denaturation. What is the evidence to support this?

a. Crystal structures show that most nonpolar side chains are buried inside a protein, which is tightly packed and which excludes water. Studies show that as the surface area of amino acid side chains increase, the free energy of transfer of amino acids from water to ethanol becomes more negative.

Figure: Transfer of amino acids from water


(Review free energies of transfer of hydrophobic groups in Lipids in Water - Thermodynamics)

b. Low temperature denaturation of proteins: It has been observed that proteins can denature at low temperatures (less than 0°C), suggesting that nonpolar residues become more "soluble" in water at low temperatures (i.e. they move from the more hydrophobic interior of a protein to the more polar outside). Compare the solubility of nonpolar gases like CO2 or N2, which are more soluble at low temperature. As you heat solutions of nonpolar gases in water, the gases become less soluble as evidenced by bubble formation (i.e. phase separation of dissolved gases as they become more insoluble). If protein behavior is governed by this same behavior (greater solubility of nonpolar groups at low temperatures), it would suggest that proteins might denature at low temperatures (leading to increased exposure to water of the nonpolar side chains). This phenomenon has been observed.

c. Protein stability affected by different salt species: Over 100 years ago, Hofmeister determined the effectiveness of different cations and anions of salts to precipitate blood serum proteins in the 0.01 - 1 M concentration ranges. The series is shown below:

Cations: NH4+ > K+ > Na+ > Li+ > Mg2+ > Ca2+ > guanidinium

Anions: SO42- > HPO42- > acetate > citrate > Cl- > NO3- > ClO3- > I- > ClO4- > SCN-

  • A salt from pairs of the first ions in these series (for example, (NH4)2SO4), when added to aqueous solutions of proteins, precipitate the native form of the protein. We must account for the fact that it precipitates the protein, and that the protein is precipitated in the native, not denatured, state. More on why it precipitates proteins in a moment. The first ion in each series increases the surface tension of water (making it harder to make a cavity in the water to fit the nonpolar molecule). This decreases the solubility of nonpolar molecules. These "salt-out" the nonpolar molecules, promoting not dissolution in water but aggregation followed by a phase separation. By analogy, they will stabilize the native state since the buried hydrophobic side chains would have a decreased propensity to move out into the aqueous environment.
  • The last ions of the series have less effect on surface tension, and hence increase the solubility of nonpolar molecules ("salt-in"). By analogy, they will destabilize the native state since the buried hydrophobic side chains would have an increased propensity to move out into the aqueous environment.
  • Hofmeister Series

Figure: Hofmeister Series


The solubility of benzene in aqueous salt solutions of this series increases from left to right, just as native protein stability decreases from left to right (i.e. the protein's nonpolar core residues become more "soluble" in water, leading to its denaturation).

Additives to proteins that increase the stability of the folded state of the protein also tend to decrease their solubilities. These additives are excluded from the preferential water hydration sphere around the protein (negative binding of these agents). Denaturants in contrast tend to increase protein solubility and interact preferentially with the protein surface. In their presence, proteins respond by increasing their surface area by denaturation. For stabilizers, proteins try to minimize their surface area by staying "native" and aggregating to form a precipitate, both of which minimizes the surface area from which the stabilizing agent is excluded.

The main effect of dissolved ions on water structure has been thought to involve changes in H bonds (either enhancers/structure makers or inhibitors/structure breakers) which correlate with salting-in or salting-out effects of various ions. Many techniques have been used to study these interactions:

  • viscosity: inferential information on structure
  • diffraction (x-rays/neutrons): gives information on coordination number of solvation shell (static information)
  • NMR: information on average relaxation of bulk and hydration sphere water around ions (dynamic information)
  • molecular dynamics simulations: gives insight into short but not long range interactions between ions and water

Recent studies have provided conflicting support for the notion of structure makers/breakers. New research (Omta et al. 2003) has used femtosecond mid-infrared pump-probe spectroscopy to study actual H-bonds between water molecules in salt solutions (Mg(ClO4)2, NaClO4, and Na2SO4). In pump-probe spectroscopy, a sample is excited with a short pulse (pump) and, after a short time lag, with another pulse (probe), which interacts with the excited state. The linear-polarized infrared pulses (pump) were used to excite OH groups in solution, followed by a probe pulse which was polarized 45 degrees compared to the pump pulse. Only those excited OH groups that had rotated in the time interval between the pump and probe would be excited by the probe. Using this technique, the time frame for reorientation of the OH groups, which is related to the "stiffness" of the H bonds, can be determined. The salts had no effect on the rotational motion of bulk water outside of the first hydration shell, which suggests that salts have no effect on the H bond networks in bulk water. Mg2+ ions are considered structure making, as the ions greatly increase the viscosity of water, brought about supposedly by increased H bonds among water molecules. This study does not support this model. Increased viscosity of Mg solutions must be attributed to those ions directly interacting with water molecules. The solution can be modeled as bulk water with small rigid spheres of ion + first hydration sphere. Clearly, much more experimental and theoretical work must be performed to gain structure insight into the role of salts on water structure. Until then, we will continue to try to understand the effects of different salts on water structure in descriptive terms and with use of thermodynamic quantities.

d. Conservation of hydrophobic core residues: These residues are highly conserved and correlated with structure.

e. Urea denatures proteins: 8M urea is often used to denature proteins. People used to think that urea competed with the intrachain H bonds and hence unraveled the protein. The arguments above with H bonds dispute this contention since water should then denature protein. How does urea denature proteins? It has been shown that the free energy of transfer of the nonpolar amino acids into 8M urea is increasingly negative as the side chains become bigger and more nonpolar.

Figure: Free energy of transfer of the nonpolar amino acids into 8 M urea


This is also true for denaturation by guanidine hydrochloride. Urea also increases the solubility of nonpolar molecules in a manner proportional to their surface area.

Apparently urea binds preferentially to the protein surface, and hence tends to increase the protein's surface area and hydrophobic exposure, and denature proteins. However, note in the figure below that glycerol, a bigger polar but uncharged molecule, stabilizes the native state. This pair of uncharged additives has correspondingly similar effects on protein stability as does the charged guanidine HCl/ammonium sulfate pair.

Figure: How reagents might interact with the surface of the protein.


Thermodynamic cycles can make it easier to visualize these transitions by breaking the denaturation and perturbant (urea, etc.) interaction into two separate steps, which, when added, lead to the final state. The next figures shows such a thermodynamic cycle for urea denaturation of proteins.

Figure: Thermodynamic cycle for urea denaturation of proteins


What accounts for this preferential binding or exclusion from the hydration sphere of a protein by these charged and uncharged additives? How do they affect ΔG0 and Keq for the processes shown above? Since denaturation is associated with increased protein size and alterations in the nature of solvent-accessible groups, it would follow that changes in the propensity for a solute (like urea or guanidine HCl) to partition into the hydration sphere (local water near the surface) of a protein compared to bulk water would change the ΔG0 or Keq for protein unfolding. A thermodynamic value, KP = mlocal/mbulk, can be used to quantitate the propensity for a solute to partition into these two phases, where m is the molal concentration of partitioning solute in the hydration sphere (local) or bulk solvent. A KP = 1 suggests no preference of solute. If KP < 1, the solute is preferentially excluded from the hydration sphere (resulting in native protein stabilization and aggregation), and if KP > 1, it is preferentially included, leading to denaturation. Since KP is an equilibrium partition coefficient, it is independent of solute concentration.

Courtenay, Capp, and Record (2001) used vapor phase osmometry to determine solution osmolality (Osm) and in turn preferential interaction coefficients (ΔGm3), from which KP values could be determined, using the formulas below, where a3 is the activity of component 3 - solute - in solution. (1 is solvent, 2 is protein, 3 is solute.) For example m1 = molality of bulk water = 55.5 mol/kg and b1 is the biopolymer hydration per square angstrom..

ΔGoobs = -RTlnKobs

-(1/RT)ΔGoobs= lnKobs

-(1/RT)( MΔGoobs/Mlna3)T,P = (MlnKobs/Mlna3) = ΔGm3 = (Mm3 /Mm2)T,P,m3 and

ΔGm3 = (m3bulk(KP -1) b1 ASA)/m1

Using BSA as a protein solute, they tested three destabilizers of native proteins: urea, guanidinium HCl, and guanidinium thiocyanate. Graphs of osmolaity vs destabilizer concentration (m3) showed almost linear increases in both the presence and absence of BSA, but with lines that crossed for the guanidium salts (lower osmolality in presence of BSA), which indicated strong protein:solute interaction.
Hence ΔOsm=Osm(+BSA)-Osm(-BSA) < 0. For protein stabilizers, such as glycerol, glycine, and betanine, ΔOsm > 0, indicating preferential exclusion of the solute from the protein hydration sphere.

Next, they determined, using vapor phase osmometry, the preferential interaction coefficients (from which KP values could be calculated) for BSA in the presence of increasing molal concentration (m3) of destabilizers. The values were always positive and increased with m3. KP values for the native protein were calculated to be 1.00 for KCl (control), 1.10 for urea, 1.30 for GuHCl, and 2.0 for GuSCN.

For urea, previous calculations by the group showed the KP for urea and denatured proteins ws the same as for native protein (1.10). Since it preferentially prefers to partition into the hydration sphere, and there is more hydration sphere in the larger denatured protein, urea drives protein unfolding.

Using the guanidinium pair, they could resolve the KP values into that for the cation (GuH+) and the anion. For BSA, KP for GuH+ = 1.60 and 2.4 for thiocyanate. Similar calculations could be made of KP for the denatured state of BSA. The values for GuHCl and GuH+ were 1.16 (compared to 1.30 for the native protein) and 1.32 (compared to 1.60 for the native protein).

Why does urea preferentially partition equally into the hydration spheres of native and denatured proteins, but GuHCl and GuH+ partitioned more into the native state? (Note however, that Kp for both N and D states were positive, leading to protein denaturation as with urea since the surface area and amount of hydration sphere greatly increases in the denatured state. )

Urea appears to partition selectively into the region near the peptide backbone and not charged or nonpolar surfaces. Theoretical and experimental work show that the a constant percentage (13%) of the surface of a protein, whether in the native or denatured state, is composed of the peptide backbone. This make intuitive sense since the backbone extends over the entire length of the protein. However, the % of the surface with charged side chains decreases in the denatured state since then number of charged and polar side chains are a small fraction of the entire polypeptide backbone. The lower distribution of GuHCl and GuH+ into the hydration shell of denatured BSA can be accounted for by different distributions of surface polar and charged groups in the native and denatured state as shown in the figure below. For example, the percentage of charged groups on the surface of native vs denatured BSA is 29% vs 4%.

It appears that GuHCl partitions into the hydration sphere near the peptide backbone and charged side chains. Urea appears to partition selectively into the region near the peptide backbone and not charged or nonpolar surfaces. This work gives a possible theoretical underpinning to the qualitative effects summarized in the Hofmeister series.

Table: Summary of Partition Coefficient (KP = mhydration sphere/mbulk) for urea and GuHCl effects on protein stability for BSA

Solute KP N KP D Select partition into % surface polar backbone N,D % surface charge N,D % DSA N↔D: backbone, charge Effect on N↔D
KCl 1 1 neither 13,13 19,9 13,4 none
urea 1.1 1.1 hydration sphere 13,13 19,9 13,4 shift to D
GuHCl 1.3 1.13 hydration sphere 13,13 19,9 13,4 shift to D



Figure: Water Accessible Surface Area in Protein Unfolding


Throughout the semester we will be discussing equilibria and how they may be shifted. The diagram below shows the first of many diagrams which will show cumulative examples of shifting equilibria.


Our understanding of hydrophobic interactions has changed dramatically in the last several years. This is not reflected in most textbooks. The hydrophobic effect mean different things to different people. Some refer to the transfer of nonpolar solvents to aqueous solution. Some refer to the same phenomena only if the effects have a unique temperature dependency. Other refer to the ordering of water around nonpolar residues. The most recent explanation centers around the unique temperature dependencies of the transfer reactions. Before we can understand it, here is an interesting bit of data. If you dissolve one mole of methane in hexane, the volume of 1 L of hexane changes 60 ml, but if done in water, the water volume changes 37 ml, indicating that water molecules pack more efficiently around nonpolar molecules then in its absence.

Let's now consider the thermodynamic aspects of the hydrophobic effect, as we did for micelle and bilayer formation. In a brief summary, we found the the free energy of transfer of an amphiphile from aqueous solution into a micelle, for example, was disfavored enthalpically (unexpectedly) but favored entropically (also unexpectedly until we included solvent in our model). These experiments were done at one temperature and gave us our first initial understanding of the hydrophobic effect. We will expand on this view by looking at the enthalpic and entropic contribution to the transfer of benzene into water as a function of temperature. This will lead us to a more modern view of the hydrophobic effect.