(Note: All previous parts in the Complexity Explained series by Dr. Vinod Wadhawan can be accessed through the ‘Related Posts’ listed below the article.)
How did life originate on Earth? Chemical or molecular evolution preceded the emergence of life. Under the influx of low-entropy energy from the Sun, and aided by the presence of certain rocks, atoms and molecules underwent chemical reactions resulting in the emergence of molecules of higher and higher information content or complexity. This article explains how this occurred.
8.1 From Atoms to Molecules
The chemical symbol H is used for an atom of hydrogen, which is the first element in the periodic table of elements. It has a nucleus, which is just a proton in this case, and there is an electron orbiting around the nucleus. The electron has a negative charge, exactly equal in magnitude to the positive charge of the proton. Taking this quantity as the unit of charge, we say that an H atom has a charge number 1 (Z = 1). Taking the mass of the proton as the unit mass, we say that H has a mass number 1 (A = 1). The electron is ~2000 times lighter than the proton.
Element number 2 in the periodic table is helium (chemical symbol He). There are two protons in its nucleus, and two electrons orbiting around the nucleus. There are also two neutrons in the nucleus. Neutrons are so called because they have no charge. The mass of a neutron is not very different from the mass of a proton. So, for the He atom, Z = 2, and A = 4.
Life on Earth is based on organic chemistry, i.e. the chemistry of the carbon atom, denoted by the symbol C. For this atom, Z = 6, and A = 12.
A molecule of hydrogen is denoted by the symbol H2. It consists of two nuclei of hydrogen, and there are two electrons orbiting around them. Why does hydrogen ‘prefer’ to exist as H2, rather than as H? Because H2 is more stable that H. Why? Consider the two electrons of H2. Quantum mechanics tells that they have no individuality, and are therefore indistinguishable. Let us consider any of them. Since positive and negative charges attract one another, the electron stays close (but not too close) to the two nuclei. [But for the Heisenberg uncertainty principle of quantum mechanics, the electrons of all the atoms would have gone right into their nuclei, and you and I would not be here, discussing chemical complexity!] Naturally, the positive charges on the two nuclei of H2 are better than only one positive charge in H, when it comes to exerting an attractive force on the electron. Thus H2 is more stable (it has a lower internal energy) than H because the former is a more strongly bound entity. Thus H atoms form H2 molecules because by doing so the overall free energy gets reduced (the second law of thermodynamics demands that the free energy be as small as possible). Formation of H2 from two atoms of H is an example of evolution of chemical complexity. More information is needed for describing the structure and function of H2, than of H.
What is the nature of the bonding between the two atoms of H2 or H-H? It is described as covalent bonding. Each of the two H atoms contributes its electron to the chemical bond between them, and the two electrons in the bonding region belong to both the nuclei.
Another kind of chemical bonding is the so-called electrovalent bonding (also called ionic bonding). It is the bonding that occurs between oppositely charged ions. Take sodium chloride (NaCl). For the Na atom, Z = 11, and for the chlorine atom, Z = 17. The laws of quantum mechanics are such that an atom of Na is more stable if it is surrounded by only 10 electrons, instead of 11. Similarly, Cl is more stable if it has 18 electrons, rather than 17. They can solve the problem together by getting readily ‘ionized’; i.e. an Na atom can become a positively charged ion Na+ by losing an electron (called the valence electron), and a Cl atom can become a negatively charged ion Cl– by gaining an electron. The two oppositely charged ions can lower the potential energy (and therefore the free energy) by coming close to each other, thus forming an ionic bond between them.
The third important and generally strong type of bonding is metallic bonding. It occurs in metals like aluminium (Al), copper (Cu), silver (Ag), gold (Au), etc. Take the case of Al. For it, Z = 13. But like an atom of Na considered above, it is more stable if it has just 10 electrons around the nucleus. So Al atoms, when in the close vicinity of one another, lose their three valence electrons to a common pool, and these valence electron become the common property of all the Al ions. A lump of Al metal is held together by this cloud of negatively charged electrons, compensating for the positive charges on the Al ions.
8.2 The Hydrogen Bond and the van der Waals Bond
The covalent, electrovalent, and metallic bonds described above are the so-called primary bonds. They are strong bonds. Diamond, for example, consists of covalently bonded carbon atoms, and is a very hard material. In metals also the atoms are strongly bonded to one another, as are the atoms in a crystal of sodium chloride in which the electrovalent interaction dominates. There are a number of other types of bonds or interactions which are substantially weaker, but are very important for biological systems in particular, and soft matter in general. Particularly ubiquitous is the hydrogen bond. Take the example of water, H2O or H-O-H. The oxygen atom forms covalent bonds with the two hydrogen atoms. Each such covalent bond (O-H) has two electrons associated with it, one coming from hydrogen and one from oxygen. The electron distribution around the two hydrogen nuclei in such a bond is not like that in a symmetrical bond like C-C in the structure of diamond. The oxygen nucleus has a charge number 8, which is much more than the charge number 1 of H, so it hogs a larger share of the electron charge cloud associated with the bond (we say the oxygen atom is very electronegative). This makes the nucleus of the hydrogen atom somewhat less shielded by the electron which was orbiting around it when there was no bonding of any kind. For similar reasons, the oxygen nucleus and its charge cloud of electrons are together a little more negative than they would be in an isolated atom of O. The end result is that the water molecule is like a little dipole. It has two positive ends and a negative end. All the water molecules are dipoles, so they tend to orient themselves such that a positive end (the hydrogen end) of one molecule points towards the negative end (the oxygen end) of another molecule. So we speak of hydrogen bonds, denoted in this example by O-H…O.
The most crucial aspect of the hydrogen bond in the evolution of chemical and biological complexity is that it is of intermediate strength, not as strong as the covalent bond, and yet not as weak as the so-called van der Waals bond (or the London dispersive bond). The van der Waals interaction is very weak, though always present between any two atoms. Quantum-mechanical fluctuations in the electronic charge cloud around an atom can result in a transient charge separation or dipole or multipole moment, and the electric field of this multipole induces a multipole moment on any neighbouring atom. This results in a small attraction between the two atoms.
The energy required to break a chemical bond is a measure of its strength. The melting point of a solid is an indicator of the strength of the weakest bonding in it. The covalent bond is the strongest, with a typical bonding energy of ~400 kilocalories (kcal). The electrovalent bond is typically half as strong as the covalent bond. The metallic bond shows a wide range of strengths, two extreme examples being the bonding in mercury on one extreme, and the bonding in tungsten on the other. The strength of a hydrogen bond is typically 14 kcal. And van der Waals bonding involves energies below 1 kcal. The most relevant fact for our purpose here is that the energy involved in hydrogen bonding is typically only ~10 times larger than the energy of thermal fluctuations, but is still much lower than the energy of a typical covalent bond. At typical temperatures at which biological systems exist, it is difficult for thermal fluctuations to break covalent bonds, but there is a fairly good chance that they can break hydrogen bonds.
8.3 The Hydrophobic and Hydrophilic Interactions
We have seen above that water is an aggregate of tiny dipoles. We say that it is a polar material. By contrast, there are a large number of ‘hydrocarbons’ which are nonpolar materials. [A hydrocarbon is a compound made predominantly of hydrogen and carbon atoms.] In contrast to the O-H bond in water, which is a bond with a dipole moment, the C-H bond in a hydrocarbon is largely nonpolar: The two electrons forming the C-H covalent bond are shared almost equally between C and H. Thus, a C-H bond hardly results in the creation of a dipole, and therefore it does not form a hydrogen bond with a water molecule. Now suppose we mix polar and nonpolar fluids. Segregation will occur. The nonpolar molecules will tend to huddle together because they cannot take part in the hydrogen bonding. They have a kind of ‘phobia’ for water molecules, and so we speak of the hydrophobic interaction. Since the hydrogen bond is of intermediate strength, the hydrophobic interaction is also of intermediate strength.
There are many types of organic compounds that are predominately of hydrocarbon structure, but have polar functional groups attached to them. Examples of this type are cholesterol, fatty acids and phospholipids. Such molecules have a nonpolar or hydrophobic end, and a polar or hydrophilic end. When put in water, they self-aggregate such that the hydrophilic ends point towards water, and the hydrophobic ends get tucked away, avoiding interfacing with water. This is why oil does not mix with water. By contrast, alcohol and water mix so readily that no stirring is needed; both are polar liquids. I forget the name of the king who said: ‘I do not care where the water flows, so long as it does not enter my wine!’
Beautiful self-assemblies like micelles, liposomes, and bilayer sheets may ensue because of the hydrophobic interaction.
8.4 Molecular Recognition and Self-Assembly
Let us go into some details of how the lowering of free energy occurs at the atomic scale. If two atoms are close to each other, they will bond together to form a molecule if the molecule has a lower free energy than that of the two separate atoms. Next, let us consider the possible bonding among molecules to form still larger assemblies (or ‘supramolecular aggregates‘). Things get more interesting now. The important concept of molecular recognition becomes operative here. Those types of molecules are likely to form assemblies which have a certain degree of complementarity. There are two types of complementarity to consider: That of lock-and-key-like shapes, and that of complementary charge distributions (remember, positive attracts negative). These complementarities, if present, enable two molecules to fit snugly into each other, thus lowering the overall potential energy, and thence the free energy. This is a more stable configuration because thermal fluctuations are less likely to knock the snugly-fitting molecules apart, and is the essence of chemical self-assembly in Nature. Self-assembly is like crystal growth, except that the end product may carry a lot more information; i.e. it is more complex.
The phenomenon of molecular complementarity was discovered by the Nobel Laureate Paul Ehrlich. As a student he was working on the newly discovered aniline dyes, which he used for staining biological cells. He found that each dye stained only a particular type of tissue or a specific species of bacteria, and not others. What happens is that the dye molecule moves around in the solution till it finds a binding site exactly fitting the pattern of atoms in one of its side chains. For stability, the complementarity of the ‘lock’ and the ‘key’ should be not only spatial, but also electrostatic; otherwise the specificity is not very strong: Not only the two shapes should be complementary, even the regions of positive excess charge on one molecule should be complementary to regions of negative excess charge on the other molecule. Here are some examples of spatial and charge complementarity in Nature:
- The complementarity between the active site of an enzyme and the substrate of the enzyme.
- The well-known ‘base-pair complementarity’ in DNA (deoxyribonucleic acid) and RNA (ribonucleic acid) strands. [I shall discuss this later.]
- Self-assembly of viruses and subcellular organelles.
- Receptors located on the surface of cells only bind a very limited number of substrates (often only one). The receptor is typically much more complicated (larger) than the substrate (hormone) that binds to it, as indicated in the accompanying sketch.
Supramolecular aggregates, normally formed under near-ambient conditions, do not involve covalent interactions usually. Instead, they are governed by weak, i.e. noncovalent or secondary, interactions (van der Waals; weak-Coulomb; hydrogen bond; hydrophobic; etc.). Because of this feature, the bonds in a supramolecular assembly at or near room temperature can get readily broken and re-formed, in a time-reversible manner, until the system has found its most stable configuration. Reversibility of bonding is a very important feature of self-assembly through molecular recognition.
Biological and other soft materials can self-assemble into a variety of shapes, and over a whole range of length scales. There is usually some amount of water present, and the most important factor mediating self-assembly is the hydrophobic interaction. Incidentally, self-assembly per se is a far more ubiquitous phenomenon than just molecular self-assembly. Some examples are: crystals; liquid crystals; bacterial colonies; beehives; ant colonies; schools of fish; weather patterns; even galaxies.
Self-assembly may be either static or dynamic. The former occurs in systems which are in local or global equilibrium, and which do not dissipate energy (e.g. crystals). Dynamic self-assembly is more relevant from the point of view of evolution of complexity, and always involves dissipation of energy. Here are some examples: oscillating and reaction-diffusion reactions; weather patterns; galaxies.
Weak interactions, with energies comparable to thermal energies, ensure that the bonds can be made and unmade reversibly, until the lowest-energy ordered configuration has been reached.
Growth of molecular crystals is an example of this. Reversibility also implies that the growing (self-assembling) system is close to equilibrium at all times.
8.5 Evolutionary Drug Designing
As a small digression, I want to mention here the use of the lock-and-key idea for designing drugs. Very often, for a drug to be effective, its molecular structure should be such that it can fit snugly into a relevant cleft in a protein molecule. In more general terms, drug activity is obtained through the binding of one molecule, i.e. the ‘ligand’, to the pocket of another, usually larger, molecule called the receptor. In their binding conformations, the molecules exhibit geometric and chemical complementarity, both of which are necessary for successful drug activity.
It can be very expensive to actually synthesise all those trial drugs and test their compatibility with the cleft in the protein molecule. Therefore, computers are used to carry out what is called ‘evolutionary computing‘. The computer code generates billions of random drug molecules, which it tests against the cleft in the protein. One such imaginary molecule may contain a site which matches one of, say, six sites on the cleft. This molecule is then ‘selected’ (it has an ‘evolutionary advantage’), and a billion variations of it are created, and tested using a suitable ‘fitness test’. And so on to the next generations of trial molecules, till the best drug shape is obtained. I shall discuss such ‘artificial evolution’ in a future article, after introducing the basics of biological evolution. As Kevin Kelly (1994) said: ‘Evolutionary breeding of drugs is the future of biotechnology.‘
8.6 The Tobacco Mosaic Virus
I consider here the example of the tobacco mosaic virus (TMV) to illustrate the hazy, perhaps nonexistent, line between life and nonlife. Any virus (including TMV) typically has an RNA core and a protein coating. It is possible to separate these two components, and purify and store them in the laboratory. At any later time the components can be mixed and incubated, and the TMV gets reconstituted by self-assembly. The reconstituted TMV thus not only comes back to ‘life’, it can even reproduce itself if placed on a tobacco leaf!
8.7 We Owe Our Lives to the Hydrogen Bond
Life and its evolution depend on the hydrogen bond. This bond is much weaker than the covalent bond, and yet strong enough to sustain self-assembled biological structures, enabling them to withstand the disintegrating influences of thermal fluctuations and other perturbations. Hydrogen bonding, and the associated hydrophobic interaction, has the right kind of strength to enable superstructures to self-assemble without the need for irreversible chemical reactions. There is a strong element of reversibility associated with these weak interactions, enabling the spontaneous making and breaking of assemblies until the lowest-free-energy configuration has been attained.
The amount of information contained in organized or complex matter is very high. This information is distributed among the shapes of the component molecules, and in the interaction patterns among them. The build up of this information involves a succession of stages: molecular recognition; self-assembly; self-organization; and chemical adaptation and evolution. We have already considered the first two. Let us now focus on self-organization.
Lehn (2002) defined self-organization as the ‘spontaneous but information-directed generation of organized functional structures in equilibrium conditions’. The necessary information (‘coding’) for self-organization is contained in the molecular-recognition and self-assembly proclivities of the component molecules. This coding also determines how the self-assembled edifice self-organizes into a functional structure in equilibrium. For a recent survey of the various types of coding for self-organization, see my book Smart Structures: Blurring the Distinction between the Living and the Nonliving (2007).
Self-organization is a far more ubiquitous phenomenon than something at just the molecular level. Here are some examples:
- A laser is a self-organized system. Under properly engineered conditions, photons spontaneously group themselves into a configuration in which they all move in phase, resulting in a powerful laser beam.
- A hurricane is a self-organized system. The steady influx of energy from the Sun draws water from the oceans, as well as drives the winds. Mild tropical winds may grow into an organized configuration of a hurricane when some critical threshold is crossed.
- A living cell is a self-organized system, which organizes itself all the time, depending on the environment.
- An economy is a self-organizing system. The demand for goods and services, as also the demand for labour, constantly organizes the economy in a spontaneous way, without any central controlling authority.
8.9 Chemical Adaptation and Evolution
Given a set of conditions, molecules in a system tend to self-organize so as to minimize the overall free energy. This is chemical adaptation. Now suppose this set of conditions changes. This is very likely, in fact inevitable, because we are dealing with an open system. A further round of self-organization must occur, governed as always by the second law of thermodynamics. This is chemical evolution. Moreover, the set of changing conditions, i.e. the changing environment experienced by the molecules, need not necessarily be that external to the set of molecules. Even internal changes in the molecular system present a changed environment to every member of the set. And molecular configurations are changing all the time. Thus, chemical adaptation and evolution occurs in an open system of molecules (including our ecosystem) all the time.
One can draw analogies with Darwinian evolution to see if ‘natural selection’ (i.e. molecular selection) and ‘survival of the fittest’ also occurs in chemical evolution. The answer is ‘yes’ because when the resources are limited, there is competition among the alternative molecular-reaction pathways, and only the fittest pathways can survive so far as consumption of precursor molecules and energy-rich molecules is concerned. Such considerations aroused special interest for explaining the origin of life-sustaining molecules. Some pioneering work in this direction was done by Melvin Calvin (1969), who introduced the idea of autocatalysis as a mechanism for molecular selection. I shall consider autocatalysis in the next article.
8.10 Concluding Remarks
The lock-and-key idea is crucial for explaining the evolution of molecules and molecular assemblies of increasing complexity in Nature. Two molecules may ordinarily interact only weakly, but a snug fitting of portions of the two molecules can lead to a much stronger degree of cohesion between them, because they ‘touch’ or attract each other at many points.
Another crucial factor for the chemical evolution of complexity is the reversibility of the non-covalent bonding between molecules; reversibility is the key to self-assembly. And once a stable self-assembly has got created (through molecular trial and error), generally there is no ‘looking back’. The overall large cohesive energy has a stabilizing effect. But there can still be a looking forward: If the environment changes, the self-assembled system can respond (i.e. adapt) by again exploiting the reversible nature of its non-covalent interactions. This is chemical evolution. As we shall see in due course, chemical evolution has led to biological evolution. And as I have said many times, evolution and emergent properties are the hallmarks of complexity.
I conclude by saying a few words about vesicles, which provide a good, even dramatic, example of self-organization in a nonliving complex system. Vesicles are spherical supramolecular assemblies separating an aqueous interior volume from the external solvent by means of lipid bilayers. They are also called liposomes, and are quite similar to micelles (see figure above). Given the right conditions, lipids can self-assemble into giant vesicles the size of biological cells. The basic driving force for their self-assembly is the hydrophobic interaction. As vividly described by Menger and Gabrielson (1995), vesicles can mimic the living cell in many ways, even though they are not living entities:
When a giant vesicle, which happens to have a smaller vesicle inside it, is exposed to octyl glucoside, the smaller vesicle can pass through the outer membrane into the external medium (‘birthing’). The resulting injury to the membrane of the host vesicle heals immediately. Addition of cholic acid, on the other hand, induces a feeding frenzy in which a vesicle grows rapidly as it consumes its smaller neighbours. After the food is gone, the giant vesicle then self-destructs (a case of ‘birth, growth, and death’). Such lifelike morphological changes were obtained by using commercially available chemicals; thus these processes should be assigned to organic chemistry, and not to biology or even biochemistry.