The word “aromatic” does not exactly sound friendly or inviting. It almost makes you wonder whether it refers to some family of molecules related to perfume. Perhaps because the name itself feels so unfamiliar, aromatic compounds can also seem rather easy to overlook. But once you get to know them, you find them everywhere in biology, playing important roles in everything from proteins and nucleic acids to coenzymes and signaling molecules. Their signature features—a flat ring structure and a pi (π) electron system spread across the entire ring—give aromatic compounds their characteristic stability and distinctive chemical properties. Thanks to these features, aromatic structures appear throughout biology, in molecules involved in energy production and redox reactions, in hormones and neurotransmitters, and even in the heme of hemoglobin. And among them are the nitrogenous bases, which lie at the very heart of nucleic acids.
If we look inside DNA or RNA, we can see that the bases are stacked neatly one above another. But they are not simply sitting there. Important chemical forces are at work between the stacked bases. A swan may look perfectly calm and graceful as it glides across the water, while underneath it is paddling furiously. In much the same way, countless interactions are taking place at the level of electrons between these seemingly quiet layers of bases. In this article, I want to explore what is really happening between the bases and why it matters. And to do that, we will once again need to go over a few basic ideas in chemistry first.
Base Stacking
The aromatic bases are sp² hybridized and therefore have flat, planar structures, with π electrons spread broadly above and below the plane of the ring rather than confined to particular bonds. These features provide excellent conditions for the bases to stack against one another, almost like flat plates piled neatly one on top of another. This arrangement is called base stacking. The planar structure created by sp² hybridization and the delocalized π-electron system both provide favorable conditions for this kind of stacking.
Base stacking is particularly important because it contributes greatly to the stability of the helical structures of RNA and DNA. If we picture RNA or DNA, hydrogen bonds pair up the bases that face each other across the molecule. But hydrogen bonding alone cannot account for the stability of the entire nucleic acid structure formed as these base pairs continue to stack one above another. Along with hydrogen bonding, base stacking in particular is known to make a major contribution to the structural stability of nucleic acids.
Intermolecular Forces: Dipoles, Hydrogen Bonding, and London Dispersion Forces
So how does base stacking stabilize these layers of bases? Surely they do not become stable simply because flat bases happen to lie close together. There are several intermolecular interactions at work between them, helping to hold them together. To understand where these forces come from, we first need to look at the kinds of forces that act between molecules.
The attractive forces that act between molecules are noncovalent interactions, which are weaker than covalent bonds. Some of the best-known examples are dipole–dipole interactions, hydrogen bonds, and London dispersion forces. Although they have different names, all of these forces are closely related to how electrons and charges are distributed within molecules and how those distributions change. Let us look at them one by one.
Dipoles and Dipole Moments
For molecules to attract one another, there first has to be something for them to be attracted to. The easiest thing to think of is electric charge. Positive and negative charges attract each other, while charges of the same sign repel each other. But even a molecule that is electrically neutral overall can have one side that is slightly more negative and another that is slightly more positive. This is because the electrons within a molecule are not always shared equally. Some atoms attract electrons more strongly than others, and the shape of a molecule can also cause electron density to be distributed more toward one side. As a result, charge can become partially separated within the molecule. When this creates two poles within a single molecule—one more positive and the other more negative—we call it a dipole. The word literally means two (di) poles.
When molecules with dipoles come close together, electrostatic attraction acts between opposite partial charges, while partial charges of the same sign repel each other. Both attraction and repulsion are present, but molecules tend to arrange themselves in ways that increase attraction while reducing repulsion. Such arrangements lower the overall energy and are therefore more stable. So when we look at how molecules with dipoles actually arrange themselves, we tend to find opposite partial charges facing one another, with attraction becoming the dominant feature of the interaction. It is not that repulsion is absent. Rather, the arrangements that minimize repulsion are the ones that remain stable, which is why attraction is usually emphasized when we talk about intermolecular interactions.
Now let us use a water molecule as an example to see how a dipole forms. A water molecule (H₂O) consists of covalent bonds between oxygen and hydrogen. Because oxygen is much more electronegative than hydrogen, it pulls the shared electron pairs more strongly toward its own nucleus. This shifts electron density toward oxygen, giving the electron-rich oxygen a partial negative charge (δ−), while the hydrogens, with relatively less electron density, carry partial positive charges (δ+). This separation of charge is not merely momentary. Even though the electrons are constantly moving, the uneven distribution persists on average. The water molecule as a whole is electrically neutral, but charge is spatially separated within it, creating a dipole.
This local separation of charge created by a bond between two atoms is called a bond dipole. A dipole moment expresses both the magnitude and direction of this charge separation as a vector, and is shown by an arrow pointing from the partially positive end toward the partially negative end. A vector is a quantity that has both magnitude and direction. So when several vectors are present, we cannot simply add them like ordinary numbers; we also have to consider the direction in which each one points. The tail of the dipole arrow lies on the less electronegative atom (δ+), where a small “+” sign is also shown. In hydrocarbons such as methane (CH₄), carbon and hydrogen differ only slightly in electronegativity, so the bonding electrons are shared fairly evenly and the C–H bonds are close to nonpolar.
Bond Dipoles and Molecular Dipoles: Molecular Polarity
As we can see in the image below, each O–H covalent bond in a water molecule has a bond dipole, and each dipole moment has a particular magnitude and direction. A bond dipole describes the local separation of charge within a particular bond, whereas a molecular dipole describes the overall polarity of the molecule as a whole. The presence of bond dipoles does not necessarily mean that the entire molecule is polar. Whether a molecule has an overall molecular dipole depends not simply on whether bond dipoles are present, but on how those bond dipoles are arranged within the structure of the molecule.
Let us take another look at the structure of water. A water molecule has a bent shape. Oxygen has six valence electrons and, in a water molecule, has two lone pairs and two bonding pairs, giving a total of four electron domains around the oxygen atom. In the conventional hybridization model, this corresponds to sp³ hybridization. With four electron domains, the ideal arrangement is tetrahedral, with an angle of about 109.5°, because this minimizes repulsion between the electron pairs. In water, however, the two lone pairs on oxygen repel more strongly than the bonding pairs do. This pushes the two O–H bonds closer together, reducing the H–O–H bond angle to 104.45° and giving the water molecule its characteristic bent, V-shaped structure.
Because of this bent geometry, the two bond dipole moments do not cancel each other out. Instead, they add together as vectors, leaving the molecule with a net dipole moment. In other words, water has a permanent dipole and is strongly polar overall. Its polarity arises not simply because O–H bond dipoles exist, but because the bent geometry prevents them from canceling each other. Carbon dioxide (CO₂) provides a useful contrast. Its C=O double bonds also have strong bond dipoles because oxygen is much more electronegative than carbon, yet the CO₂ molecule itself is nonpolar. This is because CO₂ has a linear structure with a bond angle of 180°. The two bond dipoles are equal in magnitude but point in exactly opposite directions, so they cancel each other completely, leaving the molecule with a net dipole moment of zero. The shape of the molecule is the deciding factor.
The Relationship Between Dipoles and Hydrogen Bonding
We have seen how water has a permanent dipole and how its bent structure gives it a net dipole moment, making the molecule as a whole distinctly polar. What is important is that this polarity does not remain confined within each individual molecule. When polar water molecules come close together, electrostatic attraction arises between the partial positive charge (δ+) on one molecule and the partial negative charge (δ−) on another. This is the basic form of a dipole–dipole interaction, and the particularly strong and distinctive dipole–dipole interaction found between water molecules is hydrogen bonding.
Hydrogen bonding is not a general attractive force that occurs between all molecules with dipoles. For a hydrogen bond to form, hydrogen—the smallest atom—must first be covalently bonded to a highly electronegative atom such as nitrogen (N), oxygen (O), or fluorine (F). This pulls electron density away from the hydrogen, leaving it with a strong partial positive charge (δ+). There must also be a nearby nitrogen, oxygen, or fluorine atom with a lone pair of electrons that can interact with this partially positive hydrogen. Both conditions are needed for a hydrogen bond to form.
Hydrogen bonds also have an important feature: directionality. In other words, a hydrogen bond is strongest when the donor–hydrogen bond and the acceptor are arranged close to a straight line, with the D–H···A angle approaching 180°. In this geometry, the partially positive hydrogen can interact most strongly with the acceptor. So a hydrogen bond is not simply a force that appears whenever two molecules come close together. A particular geometry is needed for the interaction to reach its greatest stability. This directionality allows hydrogen bonds to help define molecular structures with remarkable precision, which is why they play important roles in protein folding, base pairing in nucleic acids, RNA folding, and many other biological structures.
A single water molecule can donate two hydrogen bonds through its two hydrogen atoms and can accept up to two more through the two lone pairs on its oxygen atom. In other words, one water molecule can form hydrogen bonds with several neighboring water molecules at the same time. As a result, water molecules do not simply exist as isolated units in liquid water. Instead, they form an enormous, dynamic hydrogen-bond network in which bonds are constantly forming, breaking, and rearranging. Each individual hydrogen bond is weak compared with a covalent bond, but when huge numbers of them act together, they can produce a substantial collective stabilizing effect. Weak forces, it turns out, are not so weak when there are enough of them.
It is also useful to distinguish between hydrogen-bond donors and acceptors. Molecules in which hydrogen is bonded to a highly electronegative nitrogen (N), oxygen (O), or fluorine (F) atom—such as H₂O (O–H), NH₃ (N–H), and HF (F–H)—can act as hydrogen-bond donors through that bonded hydrogen. Another N, O, or F atom with a lone pair that can interact with the partial positive charge on the hydrogen acts as the hydrogen-bond acceptor—for example, a lone pair on oxygen in H₂O or on nitrogen in NH₃. The directionality of a hydrogen bond can therefore be represented as donor (D)–hydrogen (H)···acceptor (A), or D–H···A. The hydrogen bond is strongest when these are arranged close to a straight line.
When I think about just how important hydrogen bonding in water is to life, I find it almost awe-inspiring. As we have seen, each individual hydrogen bond is much weaker than a covalent bond, but the collective hydrogen-bond network formed by countless water molecules creates substantial attraction between them. The boiling point is the temperature at which molecules gain enough energy to overcome their intermolecular attractions and separate into the gas phase. In the case of water, it takes a great deal of energy to pull the molecules apart from this vast hydrogen-bond network, which is why water has the remarkably high boiling point of 100°C. Methane and ammonia, by comparison, boil at about −161°C and −33°C, respectively. This gives us some sense of just how unusually high the boiling point of water is. It is one important reason water can exist as a liquid under the everyday conditions found on Earth. If water boiled easily and existed as a gas rather than a liquid, there might have been no oceans or rivers, and cells could not have maintained the aqueous environment in which the chemical reactions essential for life take place. Life as we know it depends on liquid water, so living organisms as we know them might never have existed at all. Thinking about it this way makes me appreciate all over again just how fortunate—and how profoundly important—hydrogen bonding is for life.
London Dispersion Forces and van der Waals Forces
The dipole–dipole interactions and hydrogen bonds we have looked at so far have something in common: polarity—the slight shift of charge toward one side of a molecule—plays an important role. But what about nonpolar molecules such as carbon dioxide, where the bond dipoles cancel each other out across the molecule? Is there no attractive force at all between such molecules? Not quite. Even nonpolar molecules can attract one another slightly when they come very close together. The secret lies in the constant motion of their electrons.
The London dispersion force is a very weak intermolecular attraction and is sometimes used almost interchangeably with the term van der Waals forces. Here, “London” does not refer to the city. The force is named after the physicist Fritz London, who explained it in quantum-mechanical terms. Whereas the dipoles we have discussed arise from uneven charge distributions associated with differences in electronegativity or molecular polarity, London dispersion forces are universal attractions that can occur in all atoms and molecules. Even carbon dioxide, which is nonpolar because its bond dipoles cancel completely, experiences London dispersion forces with other carbon dioxide molecules. This is because even without permanent polarity, electron density is constantly fluctuating, allowing a dipole to appear for a fleeting moment.
If a molecule with a permanent dipole, created by an uneven distribution of electrons and therefore carrying partial charges, comes very close to another molecule—even one with no permanent dipole at all—it can induce a temporary dipole in that molecule. The electron distribution of the neighboring molecule becomes distorted by the electric field of the permanent dipole, creating a dipole for a very short time and allowing the two molecules to attract each other. In other words, the permanent dipole has induced a temporary dipole in another molecule. This is called a dipole–induced dipole interaction.
But what is even more interesting is that a permanent dipole is not needed at all. Even in two molecules with no permanent dipoles, the electrons are constantly moving, so at any given moment the electron density can shift slightly toward one side, creating an instantaneous dipole. This instantaneous dipole can distort the electron cloud of a neighboring molecule and induce another temporary dipole there. For that brief moment, an attractive force arises between the two molecules. This interaction involving instantaneous and induced dipoles is the London dispersion force. Because London dispersion forces can occur in all atoms and molecules regardless of whether they are polar, they help explain why all molecules can attract one another, even if only very slightly.
For London dispersion forces to act effectively, the distance between molecules has to be just right. The molecules must come close enough for their electron clouds to influence one another, but if they get too close, repulsion between their electron clouds—associated with the Pauli exclusion principle and electrostatic repulsion—rises sharply. London dispersion forces therefore become important at close range, where the molecules are near enough to influence one another but not so close that strong repulsion dominates. They act over short distances and decrease rapidly as the molecules move farther apart.
Taken one at a time, London dispersion forces are extremely weak, have little directionality, and appear only fleetingly as electron distributions fluctuate. Individually, they may seem almost insignificant. But here is the twist: if molecules are large or contain many electrons, and if they come very close together over a broad contact area, enormous numbers of these tiny attractions can occur at the same time. Their combined effect can then become far from negligible. The gecko, famous for its remarkable ability to cling to surfaces, gives us a striking example of how powerful van der Waals forces can become when they act collectively. Experiments have shown that geckos can cling to vertical surfaces while supporting loads many times their own body weight. The secret lies in the microscopic structure of their feet. Each foot has millions of tiny hairs called setae, and the ends of these hairs branch into hundreds of nanoscale plates called spatulae. These countless tiny structures make very close contact with surfaces such as walls or plastic, maximizing the contact area. As the saying goes, there is strength in numbers. Each individual interaction is extremely weak, but when countless van der Waals interactions act simultaneously across a large contact area, together they can produce substantial adhesion. This is closely related to the base stacking we will look at shortly.
We have now looked at several intermolecular forces. Ultimately, intermolecular attraction arises from asymmetry in the distribution of electrons. When differences in electronegativity and molecular structure produce a separation of charge that persists on average, a permanent dipole forms, providing the basis for dipole–dipole interactions and hydrogen bonding. Even in molecules without permanent dipoles, however, electrons are constantly moving, so momentary imbalances in charge can still arise. An instantaneous dipole can then induce a dipole in a neighboring molecule, giving rise to London dispersion forces. So although intermolecular forces may seem to come in several different forms, what really differs is the degree and manner in which electrons become unevenly distributed. At their root, they all begin with asymmetry in electron density—an imbalance in the distribution of electrons.
We have taken a rather long journey through chemistry as a stepping stone toward understanding base stacking. Now that the groundwork is done, let us get to the heart of the matter. How do the forces we have just explored actually work between bases stacked neatly one above another?
Base Stacking and van der Waals Forces
Aromatic molecules have broad electron clouds extending above and below their planar rings. These clouds are formed by π electrons that are not confined to particular atoms but are spread across the ring as a whole. These widely distributed π electrons are also constantly fluctuating, so their distribution shifts slightly from moment to moment. At any given instant, electron density may become somewhat greater in one region of the aromatic ring and lower in another.
When electron density becomes temporarily uneven in this way, a region with relatively more electrons takes on a fleeting partial negative charge (δ−), while a region with lower electron density takes on a partial positive charge (δ+). Unlike a permanent dipole produced by differences in electronegativity, this is an instantaneous dipole created by momentary fluctuations in the π-electron cloud. This instantaneous dipole can then distort the electron cloud of a neighboring aromatic base and induce another dipole there. A weak attractive force arises between the instantaneous dipole and the induced dipole, and this is the London dispersion force.
And just as we saw with the gecko, an individual London dispersion interaction may be extremely weak, but in base stacking, where broad surfaces come into very close contact, these tiny effects can add up to something that is far from negligible. Once again, there is strength in numbers. When large numbers of π electrons participate in these interactions at the same time, dispersion forces can become one of the important forces that help stabilize base stacking.
Quadrupoles and the Geometry of Base Stacking
The image below uses benzene, the simplest aromatic molecule, to help explain base stacking. Benzene’s π-electron cloud is concentrated above and below the ring, making these regions relatively electron-rich, while the hydrogens in the C–H bonds around the edge of the ring are relatively electron-poor and have a partially positive character. Benzene is electrically neutral overall and has no permanent dipole, but its charge is still distributed through space in a characteristic pattern. This kind of charge distribution can be described as a quadrupole, a distribution more complex than a dipole and, literally, one involving four poles.
When two benzene rings come together, these electrostatic features act alongside the London dispersion forces we have already discussed and the electron-cloud repulsion that rises sharply at very short distances. As a result, rather than stacking perfectly face-to-face like two slices in a sandwich, the rings can adopt stable arrangements such as a slightly offset parallel-displaced geometry or a T-shaped geometry. Parallel-displaced stacking, which is commonly observed in molecular structures, allows the π-electron clouds to overlap partially and gain attractive dispersion interactions while reducing the repulsion that would arise if the electron clouds came too close. The bases in DNA are likewise stacked in offset arrangements rather than lying perfectly on top of one another.
DNA Twist and Base Stacking
If we look at B-DNA, the most common form of DNA under physiological conditions, the base pairs are not stacked perfectly straight one above another. Instead, each base pair is rotated by about 36° relative to the one above it. As this rotation repeats along the molecule, DNA takes on its familiar spiral-staircase shape. This rotation is called the helical twist. The helical structure emerges from the combined effects of base stacking, structural constraints imposed by the sugar–phosphate backbone, and other interactions. Because of this rotation, neighboring bases do not overlap perfectly but stack in an offset arrangement.
The two bases within a single base pair are also slightly twisted relative to each other, somewhat like the blades of a helicopter propeller. In other words, one base tilts slightly upward while the other tilts slightly downward. This out-of-plane twisting of the two bases relative to one another is called propeller twist. Together, propeller twist and the rotation and offset between neighboring base pairs create the three-dimensional base-stacking geometry seen in real DNA. The bases are also stacked about 3.3–3.4 Å apart, a spacing that fits well with the distance at which attractive forces, including dispersion forces, are balanced against the repulsion that increases at very short range.
Base Stacking and DNA Structural Stability
To compare how much hydrogen bonding between base pairs and base stacking contribute to the stability of the DNA double helix, researchers measured the heat involved when DNA melts. These thermodynamic analyses found that base stacking contributes about 60% of the total free energy involved in stabilizing the DNA double helix, and that most of the heat absorbed as DNA melts is associated with the disruption of stacked bases. [1] This once again shows that base stacking is not simply a side effect of bases happening to pile up one above another. It plays a major role in holding the DNA double helix together and keeping it stable.
Hydrogen bonds are local interactions with clearly defined positions and directions: two form between A and T, and three between G and C. Base stacking, by contrast, is not an interaction confined to a particular atom or bonding site. Instead, several interactions act together across the broad surfaces of neighboring stacked bases. So when the DNA double helix opens, it is not only the hydrogen bonds that are disrupted. The bases also move apart and lose their orderly alignment, causing the stacking interactions to break down at the same time. During DNA replication, when helicase unwinds the double helix, both base pairing and the stacked arrangement are disrupted as the two strands separate. Opening DNA therefore requires more energy than simply breaking the hydrogen bonds between its bases.
The Hydrophobic Environment Created by Base Stacking
Base stacking contributes to the stability of RNA and DNA in another important way. As the bases stack one above another, the surfaces where neighboring bases come into contact form a relatively hydrophobic environment that water cannot easily enter. Very little space remains for water molecules to sit stably between the stacked base planes. The bases themselves do contain polar functional groups, but these are located mainly around the edges of the base planes. The interface formed when the bases stack, on the other hand, consists largely of contact between their relatively hydrophobic aromatic surfaces.
Water tends to favor arrangements in which it can maintain its hydrogen-bond network, but a water molecule trapped in the narrow space between two bases has very limited freedom in the angles and directions available for hydrogen bonding. This makes the space between the bases an unfavorable place for building that network. It is therefore more favorable for water to leave the space between the bases and return to the surrounding aqueous environment. We can think of this as similar to what happens when a protein folds and hydrophobic amino acid residues gather together away from water and become buried inside the protein. In nucleic acids as well, it is more favorable for the relatively hydrophobic aromatic surfaces of the bases to stack against one another rather than remain separately exposed to water, thereby reducing the surface area in contact with water. Put simply, the bases come together and stack away from water, and that stacked state in turn helps stabilize the nucleic acid structure. In other words, the aqueous environment itself helps stabilize the stacked arrangement of the bases.
As water molecules around the bases return to the surrounding aqueous environment, they gain greater freedom of movement and can once again participate in the hydrogen-bond network of water, which can also be thermodynamically favorable. As a result, this hydrophobic effect helps stabilize the stacked arrangement of the bases, providing another important contribution to the structural stability of RNA and DNA in an aqueous environment. If hydrogen bonding between bases contributes to recognition and selectivity between base pairs, base stacking and the hydrophobic effect associated with it provide important physical support for the overall structures of RNA and DNA.
Hydrophilic vs. Hydrophobic
Life operates in an environment of water, and most chemical reactions in living organisms take place in an aqueous environment. So it is hardly an exaggeration to say that how a molecule interacts with water can help determine its function and even its fate. Hydrophilicity and hydrophobicity are relative concepts defined in relation to water. A molecule that can form favorable interactions with water, such as hydrogen bonds or electrostatic interactions, is described as hydrophilic, whereas one whose interactions with water are relatively unfavorable is described as hydrophobic. In other words, these are not simply absolute properties that a molecule possesses on its own. They describe the molecule’s relationship with water.
Water molecules form a stable hydrogen-bond network through strong hydrogen bonding with one another. Molecules that interact well with water can remain dispersed and stable in solution, while molecules that do not interact as favorably with water may find it more favorable to gather together and reduce the amount of surface exposed to water. So when hydrophobic molecules cluster together, it is less a case of them liking one another and pulling themselves together than of being pushed out because water does not like them, so to speak. In other words, they come together not because there is some unusually strong attraction between hydrophobic molecules themselves, but because the surrounding aqueous environment makes that arrangement more favorable.
This hydrophobic effect can become an extremely important driving force. Molecules that do not interact well with water tend to hide away from it on the inside, while water-friendly molecules happily(?) remain exposed at surfaces where they can stay close to water. This kind of arrangement can strongly influence how molecules organize themselves in three-dimensional space. One of the major factors that determines how a protein folds is the hydrophobic or hydrophilic nature of its amino acid residues. As residues that avoid water and residues that prefer contact with water spontaneously arrange themselves, their distribution becomes a major factor in shaping the protein’s structure.
Cell membranes provide another example of a structure that forms spontaneously through the combined hydrophobic and hydrophilic properties of phospholipids. Their hydrophobic tails point inward, while their hydrophilic heads face outward toward the surrounding aqueous environment. Hormones also interact with cells differently depending on whether they are hydrophobic or hydrophilic. Hydrophobic hormones that can cross the membrane directly can enter the cell and interact with intracellular receptors, whereas hydrophilic hormones generally need receptors on the cell surface to transmit their signals into the cell. In biochemistry, then, a molecule’s relationship with water matters enormously. And because water makes up a large part of our bodies, hydrophilicity and hydrophobicity are important considerations throughout biological function. RNA and DNA are no exception. Their outer sugar–phosphate backbones are strongly hydrophilic, while on the inside, the relatively hydrophobic surfaces of the bases stack against one another, reducing the amount of surface exposed to water. This arrangement helps RNA and DNA maintain stable structures in an aqueous environment.
I have explained hydrophobicity only briefly here, but given the importance of water in biological systems, a more accurate and deeper explanation would eventually have to bring in the concepts of free energy and entropy. Since the purpose of this article is to better understand RNA and DNA, however, I will stop at this level for now. I have discussed in more detail elsewhere how free energy and entropy connect the hydrophobic effect with the structural stability of biomolecules, in an article on protein folding and thermodynamics, so if you are interested, you may want to come back to that later.
[References]
[1] Forces maintaining the DNA double helixhttps://doi.org/10.1007/s00249-020-01437-w
Biochemistry, 9th Edition by Mary K. Campbell, Shawn O. Farrell, and Owen M. McDougal (2018)
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2-2. 209 Polar Covalent Bonds in a Water Molecule — CC BY 3.0
2-3. Water molecule – structure and dipole moment — CC BY-SA 4.0
Image 3
3-1. 210 Hydrogen Bonds Between Water Molecules-01 — CC BY 3.0
Image 4
4-1. π⁺–π⁺ stacking of imidazolium cations enhances molecular layering of room temperature ionic liquids at their interfaces — CC BY 3.0
4-2. Molecular Motions in Functional Self-Assembled Nanostructures — CC BY 3.0
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Image 5
5-1. DNA Orbit Animated Clean — CC BY-SA 4.0
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 — Hydrogen Bonding and Base Stacking){kind=link}