Structurally, RNA and DNA are very similar. Both are nucleic acids made of long chains of nucleotides, with nitrogenous bases attached to the sugar–phosphate backbone that forms the basic framework of the molecule. But when it comes to their roles and functions, RNA and DNA are clearly very different. In the previous long article centered on RNA, we took a close look at the sugars used by these two nucleic acids and how that difference affects their stability and function. The ribose sugar in RNA has an OH group at the 2′ carbon, whereas DNA has a hydrogen instead of an OH group at the same position. The oxygen is missing. RNA is called ribonucleic acid because it contains ribose. DNA, on the other hand, contains deoxyribose, which has the same basic structure as ribose but one less oxygen at the 2′ carbon—hence the name deoxyribonucleic acid, meaning a nucleic acid containing deoxyribose. Of course, this small difference did much more than simply give the two molecules their names. Whether or not an OH group is present at the 2′ position has a major effect on the stability and reactivity of the two nucleic acids and ultimately helped steer RNA and DNA toward different roles. You could say that this tiny OH group ended up changing the fate of the two nucleic acids.
But there is another notable chemical difference between these two nucleic acids. One of the bases they use is different. RNA uses four bases—A, G, C, and U—whereas DNA uses A, G, C, and T. RNA uses U, while DNA uses T in its place. So perhaps it is time to turn our attention to the bases themselves—the nitrogenous bases that actually make up the content of genetic information. Among the countless organic molecules that exist, why are these particular bases so well suited to serve as the information units of nucleic acids? What chemical properties allow them to hold genetic information safely within their structures? And what exactly are these molecules? Let us take a closer look, one question at a time.
The Two Types of Nitrogenous Bases Used by Nucleic Acids: Purines and Pyrimidines
In biochemistry, a nitrogenous base can be described as an aromatic heterocyclic organic compound containing nitrogen atoms, in which the lone pair of electrons on some of the ring nitrogens can accept a proton and give the molecule basic properties. In other words, they are called “nitrogenous bases” because they contain nitrogen and can exhibit basicity. But the term nitrogenous base is not generally used simply as a broad name for any “nitrogen-containing basic molecule.” Instead, it is most commonly used as a specific chemical and functional term for the base component of nucleic acids—RNA and DNA. From the perspective of nucleic acid biochemistry, the standard nitrogenous bases found in nucleic acids belong to just two structural families: purines, with two fused rings, and pyrimidines, with a single ring. Of course, living organisms contain many other nitrogen-containing aromatic ring compounds, including imidazole and pyridine, but when it comes to nucleic acids, only these two families were selected. Imidazole, for example, forms the side chain of the amino acid histidine and is itself a stable aromatic heterocycle with a variety of biological functions. So this naturally raises another question: among the many nitrogen-containing aromatic rings, why do nucleic acids use purines and pyrimidines?
Why Aromatic Ring Compounds?
To understand why aromatic rings are so well suited to serve as the information units of nucleic acids, let us first look at what aromatic ring compounds actually are. Organic compounds are built around carbon skeletons, which is why the term organic compound is often used almost interchangeably with carbon compound. Carbon compounds can be classified in many different ways. If we classify them simply by the shape of their carbon skeletons, one of the most basic divisions is between open-chain compounds, with straight or branched skeletons, and cyclic compounds, which form closed rings. Cyclic carbon compounds can then be divided into aromatic and non-aromatic compounds. Aromatic compounds have a distinctive π-electron system in which electrons are spread, or delocalized, over the entire ring. This gives them characteristic planarity and high stability. Organic compounds are therefore also sometimes divided more broadly, regardless of the shape of the carbon skeleton, into aromatic compounds and all the remaining non-aromatic compounds, collectively called aliphatic compounds.
For living organisms to store and transmit information in an aqueous environment, they would have needed molecules that could remain chemically stable in water over long periods while also forming regular, repeating structures. In both RNA and DNA, the bases are stacked one above another, and the planarity and relatively rigid structure of aromatic rings are very useful properties for building this kind of arrangement. Most open-chain carbon compounds contain single bonds around which free rotation can occur, making their structures potentially too flexible to form such regular stacked structures. Typical aromatic rings such as benzene and the nucleic acid bases are different. The atoms in these rings are generally sp² hybridized, giving the rings a flat, planar structure with relatively fixed bond lengths and angles. This provides geometric stability, which is particularly well suited to something that must serve as a repeatable unit of information. In that sense, the structural properties of aromatic rings seem almost tailor-made for the information units of nucleic acids.
The Key to Understanding Typical Aromatic Rings: sp² Hybridization
Above all, one of the most important characteristics of aromatic carbon compounds is their stability. To understand where that stability comes from, we first need to understand sp² hybridization. We already looked in detail at how carbon undergoes sp² hybridization in a separate article on hybrid orbital theory. But because a basic understanding of hybridization is very helpful for properly understanding nitrogenous bases, let us briefly go over it again here. Carbon hybridizes its orbitals to optimize bonding with different bonding partners. In the conventional model used to explain hybridization, one electron is first promoted so that the s and p orbitals in the second electron shell, or L shell, which have different energies in the stable ground state, can be treated as being at the same energy level in an excited state. From this excited-state configuration, carbon can undergo different types of hybridization depending on its bonding environment. One thing is worth keeping in mind, however. This idea of electron promotion is a model that makes hybridization easier and more intuitive to understand. In an actual molecule, electron promotion does not happen first, followed step by step by hybridization and then bond formation. Rather, orbital reorganization occurs together with bond formation.
The sp² hybridization shown in the middle of the figure below illustrates how one 2s orbital and two 2p orbitals hybridize to form three sp² orbitals, which then form sigma (σ) bonds with neighboring atoms. Meanwhile, one p orbital remains unhybridized and can overlap with a p orbital on a neighboring atom to form a pi (π) bond. In a sigma bond, the orbitals overlap strongly along the axis connecting the nuclei of the two atoms. A pi bond is different: parallel p orbitals overlap above and below, or to the sides of, the atomic nuclei, with electron density distributed in those regions. The three sp² orbitals are arranged in a trigonal geometry with angles of about 120° between them, all in the same plane. Because negatively charged electrons repel one another, electron pairs tend to arrange themselves as far apart as possible to minimize that repulsion, and the number of electron pairs determines their geometric arrangement. In the case of sp² hybridization, arranging the electron pairs in a trigonal planar geometry provides the optimal spatial arrangement for minimizing electrostatic repulsion between them. We also saw that when a π bond is added to one of the σ bonds, the two together form a double bond. This type of sp² hybridization is commonly found in many organic compounds.
Using ethylene, one of the simplest molecules with sp²-hybridized carbon atoms, as an example, I have illustrated how sigma and pi bonds are formed. Ethylene is also well known as a signaling molecule that regulates fruit ripening in plants. When a pi bond is added to a sigma bond, together they form a double bond.
What Makes sp² Hybridization in Aromatic Compounds Different: The Delocalized π System
As a first step toward understanding the stability that is such an important feature of aromatic compounds, we have looked at carbon hybridization, particularly sp² hybridization. At the orbital level, sp² hybridization itself is not fundamentally different in an aromatic compound from ordinary sp² hybridization. But there is one crucial difference. In an aromatic ring, the p orbitals left behind after hybridization overlap continuously around the entire ring, creating a delocalized distribution of π electrons.
Let us look at this more closely using benzene, the simplest and most basic aromatic molecule, as an example. Benzene has the molecular formula C6H6. All six carbon atoms are sp² hybridized and arranged in the same plane to form a ring, with bond angles of about 120°. Each carbon forms three sigma (σ) bonds—two with neighboring carbon atoms and one with a hydrogen atom—leaving one p orbital that does not participate in hybridization. These six remaining p orbitals are aligned parallel to one another and overlap continuously around the entire ring. As a result, the six π electrons are not confined to any particular carbon–carbon bond. Instead, they are distributed over the ring as a whole. This is what we mean by the delocalization of π electrons.
In ethylene, which we looked at earlier, the p orbitals of the two carbon atoms overlap directly and form a π bond confined to those two atoms. This localized π bond can therefore be broken or participate in chemical reactions relatively readily. Aromatic compounds such as benzene are different. Their p orbitals overlap collectively around the entire ring, forming one continuous π-electron system. What matters here is that the π electrons do not belong to one particular bond but are spread over the ring as a whole. This delocalization of π electrons contributes greatly to the distinctive stability of aromatic compounds. The flat shape of aromatic rings and their broadly distributed electron systems also provide an important structural basis for biochemical properties such as base stacking, in which the nitrogenous bases of nucleic acids stack one above another. This is also why benzene is often drawn as a hexagon with a circle inside: the circle is a visual representation of the π-electron system spread across the entire ring.
What is worth noticing is that this collective π-electron system lies at the heart of the special stability of aromatic compounds. Because the π electrons are spread over the entire ring, the energy of the molecule is substantially lowered. As a result, aromatic compounds show unusual stability toward reactions that would disrupt the π-electron system of the ring. This high chemical stability helps nucleic acid bases maintain stable, planar ring structures. Once we see this, it begins to make sense why the bases of nucleic acids—molecules that must preserve genetic information safely—have aromatic structures.
Carbon Aromatic Rings vs. Nitrogen-Containing Aromatic Rings
But then another question comes to mind. Why did nucleic acids come to use nitrogen-containing heteroaromatic rings rather than carbon-only aromatic rings such as benzene? Benzene consists only of carbon and hydrogen. Its carbon atoms share the same sp²-hybridized environment, its electron density is highly symmetrical, and the molecule has an almost perfectly planar and stable structure. But this very symmetry and uniformity can also become a disadvantage.
In benzene, all of the carbon atoms are connected by C–C bonds with the same geometry, and the electron density is distributed very symmetrically across the entire ring. You could almost say that, chemically, benzene looks the same whichever way you approach it. As an analogy, if another molecule wants to interact with benzene, there are relatively few distinct chemical signposts telling it which direction to approach from or where to interact. When molecules interact, the ability to distinguish a particular molecule and interact with it in a particular way can be thought of as a form of information recognition. Benzene itself is so structurally and electronically symmetrical that it offers very few distinctive chemical features that could serve as such signposts.
But once nitrogen is added to the ring, the situation changes. Nitrogen is more electronegative than carbon, so it pulls electrons toward itself and shifts the electron density locally. Partial charges (δ+ and δ−) appear, and the previously symmetrical electron distribution becomes somewhat uneven. In other words, the ring now has chemical directionality. Depending on its position and bonding state within the ring, nitrogen can also use a lone pair to act as a hydrogen-bond acceptor, or, when bonded to hydrogen, act as a hydrogen-bond donor. Add functional groups such as oxygen-containing groups or amino groups at different positions, and each base develops its own pattern of hydrogen-bond donors and acceptors. Now selective interactions between molecules become possible. In the nitrogenous base adenine (A), for example, whether a nitrogen participates as a hydrogen-bond donor or acceptor depends on where it is located and how it is bonded. This is exactly the kind of chemical differentiation that allows nitrogen-containing aromatic molecules to serve as information units. By using heteroaromatic rings containing nitrogen, nucleic acid bases can keep the stability and planarity of aromatic structures while placing hydrogen-bond donors and acceptors at different positions. Each base can therefore present its own distinct chemical surface, making selective complementarity between bases possible.
Then why is the ability to form hydrogen bonds so important? Hydrogen bonds arise from differences in electronegativity. They are not nearly as strong as covalent bonds, but their directionality and selectivity are far from negligible. A single hydrogen bond may be weak, but several acting together can maintain a sufficiently stable interaction. And because hydrogen bonds form most favorably when conditions such as distance and angle are just right, they play an important role in selectively recognizing matching base pairs. More importantly, during DNA replication or the transcription of DNA into RNA, the base pairs must temporarily separate. If the bases were held together by strong covalent bonds, would that not make the whole process almost impossible? The backbone of a nucleic acid is held firmly together by covalent bonds, while the nitrogenous bases on the inside—the parts involved in information recognition—interact through noncovalent forces such as hydrogen bonds. It seems like a remarkably clever solution: information can be stored securely, yet still remain accessible for copying and modification. For much the same reason, RNA, whose flexibility is one of its strengths, can fold as bases within the molecule attract one another through hydrogen bonding. Because these hydrogen bonds can change depending on the surrounding environment, RNA can shift among a variety of secondary and tertiary structures.
In the next article, we will take a closer look at base stacking, an even more important factor in stabilizing nucleic acid structure, alongside hydrogen bonding between bases.
[Image sources]
2. Chemistry LibreTexts (CK-12 Foundation). Hybrid Orbitals – sp and sp². CC BY-NC 3.0.
3. Chemistry LibreTexts (CK-12 Foundation). Hybrid Orbitals – sp and sp². CC BY-NC 3.0.
5.Vladsinger. Benzene Representations. Wikimedia Commons.
https://commons.wikimedia.org/wiki/File:Benzene_Representations.svg, CC BY-SA 3.0
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