Chemistry

Molecules

The galactic epoch has then led to a considerable chemical enrichment of the universe through the nuclear route. From the 75/25 split between Hydrogen and Helium after the primordial nucleosynthesis (plus some traces amounts of heavier nuclei), we are now at 74/24. "Only" 2% of the universe’s ordinary matter mass is made of heavier elements, but we’ve already seen what a big impact they had in galaxy formation, stellar lifespan, etc. Half of that amount is in oxygen (1%), a quarter in Carbon (0.5%), and almost all the rest in neon, iron, nitrogen, silicon, magnesium and sulfur (0.4%). The remaining 0.1% is split between trace amounts of all naturally occuring elements, including uranium, the heaviest primordial element (0.2 parts per billion), which can then decay into even heavier ones: neptunium and plutonium.

But that is not the only way that chemical enrichment occurred. Previously, the name "molecular cloud" only referred to one molecule, really: dihydrogen \(\text{H\_2}\). In a universe with only hydrogen and helium, the chemical bond between two hydrogen atoms is the only one at hand ("dihelium" only exists at extremely low temperatures and even then is highly unstable).

Molecular structures are the products of intramolecular forces, which are derivative of electromagnetic interactions. We can draw an analogy with the residual nuclear force: it keeps protons and neutrons together in a nucleus, despite the nucleons already being color-neutral. But the strong interaction is intense enough that quarks from distinct nucleons can also exchange gluons and keep their respective containers tightly knit. Similarly, at sufficiently low temperatures, a plasma state can turn into an atomic gas when electrons get captured by nuclei and form atoms: electrically neutral objects. When the ambient agitation gets even weaker, these atoms can also become bound to each other.

The key driver for intramolecular forces is the specific nature of atoms’ electronic configuration. Being quantum particles, bound electrons in an atom don’t have an infinity of states they can be in. There is a finite number that is entirely predictable by the theory. It is a consequence of the fermionic nature of electrons: following Pauli’s exclusion principle, no more than one fermion can occupy each state. Every electron is characterized by 4 quantum numbers:

The principal number \(n\) determines the "shell".
The angular momentum number \(l\) runs from \(0\) to \(n-1\) and determines the "sub-shell".
The magnetic number \(m\) runs from \(-l\) to \(+l\).
The spin \(s\) takes value \(-1/2\) and \(+1/2\).

We then have \(2n^2\) possible electrons per shell \(n\).

Now, what would it mean to talk about the electric configuration of oxygen or of gold? We’ve seen that elements are characterized by their atomic number \(Z\), the number of protons, meaning the nucleus’s electric charge is \(+Z\). Different numbers of neutrons lead to different isotopes, but it’s still the same element. That’s a matter of nuclear physics. Different electrons around that nucleus and it’s still the same element, but the difference is now a matter of chemistry. For neutral atoms, the electronic orbitals contain \(Z\) electrons. In all other configurations, we talk of an ion instead.

As we go down the list of neutral atoms from hydrogen (\(Z=1\)) onwards, new electrons are then added. We’ve seen that each one is characterized by a quadruplet \((n,l,m,s)\). These quantum eigenstates have distinct energy levels, so if an atom has \(Z\) electrons, they cannot just be in any set of states: they always tend to occupy the lowest-energy states. Indeed, if they didn’t (which is always possible, probabilistically speaking), the expected half-life of that configuration would be vanishingly short, as the offending electron would soon jump down to the available lower energy level, emitting a photon in the process.

Quantum physics can help us determine the order in which the electron states are occupied, and thus to know exactly which "spots" are occupied by the electron of given atom. That information is called the electronic configuration of an element.

The start of the sequence is simple enough: the first shell is filled first. It is called \(1\text{s}\): \(1\) because it is the first shell; \(\text{s}\) because it denotes the first subshell of any shell. \(1\text{s}\) has space for two electrons, with opposite spins but same energy levels. Hydrogen has 1 electron, so the electron configuration is written \(1\text{s}^1\). For helium, we write \(1\text{s}^2\). After that, we move on to the second shell, which has two sub-shells: \(2\text{s}\) and \(p\text{s}\), with respectively 2 and 6 electrons. We do the same for the third shell, filling out its first two sub-shells. We reach Argon, whose electronic configuration is: \(1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^12 3\text{p}^6\).

We might expect the third sub-shell, \(3\text{d}\), to be filled next, but that is not the case. Instead, we move to \(4\text{s}\), before coming back to that one. That is because electron energy increases with \((n+l)\). The \(4\text{s}\) subshell is defined by \((n=4,l=0)\), so it had to come before \(3\text{d}\) \((n=3,l=2)\). Between two subshells with same \((n+l)\), the lower shell comes first, which is why \(3\text{p }\) \((n=3,l=1)\) came before \(4\text{s}\). That is the Aufbau principle.

What is the sequence within a given subshell? Hund’s rule states that the first \(2l+1\) electrons will fill states with distinct \(m\), in no particular order since they have the same energy; but the take-away is that no opposite-spin pair of electrons sharing the same \((n,l,m)\) numbers will be formed within the subshell until there is no option.

The principle seems to be that subshells are fully filled, one after another. That is not entirely true, and irregularities are due to the complexity of electronic clouds’ self-interactions. For instance, Nickel (\(\text{Ni}\)) is in this configuration: \([\text{Ar}] 4\text{s}^2 3\text{d}^8\) (where \([\text{Ar}]\) is a short hand that means: the whole of Argon’s electronic configuration, i.e. all sub-shells filled up to \(3\text{p}\)). We would expect the next element to add an electron to the outermost subshell \(3\text{d}\). But not, copper () is in this configuration: \([\text{Ar}] 4\text{s}^1 3\text{d}^{10}\). It "deprives" the fourth shell of one of its electrons. With the next element, zinc (), we come back to \([\text{Ar}] 4\text{s}^2 3\text{d}^{10}\).

A key distinction we’ll operate now is between core electrons and valence electrons. Core electrons are "locked in", and unless the medium is fully turned into a plasma, they don’t interact with the outside world at all. Valence electrons can interact with the outside world, and they do: they can elope with another atom, form bonds, or get kicked out! The rule is simple: electrons on any subshell that is not full are valence electrons; electrons on the outermost shells are valence electrons. Oftenwise these two sets are the same. But not always, as we’ve seen. Nickel has 10 valence electrons, 8 from the unfinished \(3\text{d}^8\) subshell, and 2 from the finished, but outermost \(4\text{s}^2\). The rest of the electrons form the core shell and can be safely ignored. In fact, the valence electrons only "see" the nucleus through the core shell. If there are \(v\) valence electrons, they tend to see the [core shell + nucleus] ensemble as if it was a nucleus with charge \(+v\). That is why there is a strong kinship between elements sharing the same valence electron configuration.

This is reflected in the organization of the periodic table. A new row (period) starts when a new shell gets its first electrons. Each column (group) includes elements with the same valence shell configuration.

Now, all intramolecular (or inter-atomic) forces are based on the electron configurations of atoms. Every atom can become a cation (a positive ion) by losing an electron: that requires energy, called the ionization energy. Every atom can also become an anion (a negative ion) by attracting and holding on to an extra electron. The energy that is then released is then called the electron affinity (a negative electron affinity would then mean it would take energy to add an electron).

The following trends can be identified:

Going down a group, atoms have the same electron configuration, but on farther and farther shells. They also become more and more prone to lose electrons, all else being equal. That is because electrons on more distant shells are subject to a smaller attrative force from the nucleus.
Going across a period, we are successively filling up a number of subshells. An atom with fully-filled shells is extremely stable (only helium and neon are in this case); an atom with only fully-filled subshells is also very stable. If a neutral atom is just missing one electron to complete a sub-shell that already has seven, then its electron affinity and ionization energy will be high: it is "inclined" to absorb an electron and be in a stable electronic configuration, and "reluctant" to let an electron go, getting further away from that stable state. Conversely, an atom with only one electron in the outermost sub-shell has low electron affinity and ionization energy: it should be cheap to knock out that extra electron, and unlikely to add another one to the mix.

So, the more towards the top-right corner of the periodic table we go, the more atoms tend to attract, rather than lose, an extra electron. This is captured under the notion of chemical metallicity (which is distinct from the astrophysical concept, that only labels hydrogen and helium as nonmetals). In chemistry, nonmetals include hydrogen, carbon, nitrogen, oxygen, phosphorus, chlorine, selenium, bromine, iodine, and all the noble gases (elements from the right-hand side group with a \(n\text{s}^2n\text{p}^6\) outer shell): helium, neon, argon, krypton, xenon and radon. A buffer group called quasimetals or metalloids is less well-defined, and includes atoms like boron, silicon and arsenic. All other elements are metals.

The metallicity of elements directly impacts the type of chemical bonds that they can form.

Noble gases have very stable electronic configuration, and hardly ever bond spontaneously with other elements. They are considered chemically inert.
Other nonmetal elements tend to form covalent bonds between themselves. Remember that they have a tendency to seek extra electrons to complete their outermost subshell. A covalent bond achieves that through the sharing of valence electrons. The simplest case is the dihydrogen molecule: each hydrogen atom only has one electron, but a full subshell requires two. Be that as it may, it cannot be expected, in a neutral medium, to find a wandering free electron. In a dihydrogen molecule, the orbitals overlap in such a way that each hydrogen atom "feels" that it has a full subshell. As we fill a full subshell, we pair electrons of opposite spins together. Although individual electrons have no identity, we can represent this sharing as the "pairing" of electrons from distinct atoms. Sometimes, two or three of electrons will be shared between two atoms, leading to double or triple bonds (or much more rarely, quadruple, etc.).
Sometimes, covalent bonds can only exist in the form of a crystal network, with a repeatable unit. The typical example is diamond, a carbon crystal. Objects that can be assembled from covalent bonds and exist independently are called molecules.
Due to the quantum constraint on covalent bonds, molecules have a specific geometry, and therefore polarity.Those features are crucial, not only to chemical reactions themselves but to the emergence of complexity at higher levels.
Metals and nonmetals, when they are in each other’s vicinity, tend to develop ionic bonds. Metals seek to lose atoms, and nonmetals to gain them. The energetic path of least resistance is then for the metal to become a cation, and the nonmetal an anion, with the former transferring an electron to the latter. This can be seen as an asymptotic limit to the covalent bond. In covalent bond between different elements, the relative electric charges from their nuclei generates an electric polarity, and the shared electrons are closer to one of the atoms. We could consider that in the ionic "bond", all involved electrons are so close to the anion that they are not connected to the cation anymore at all.
Unlike molecules, it is not possible to identify a single unit here. Ions are not connected to each other and can freely wander in a liquid or gaseous state. Alternately they can compress into a crystal where the same cation/anion alternating pattern is repeated (salt sodium chloride is an example).
Metals all want to lose electrons. Metallic bonding then adopts the opposite stance to covalent bonding: instead of sharing electrons and bring them into the nice cocoon of shared orbitals, they delocalize them, letting them wander freely. In a way, every electron is also shared between every single atom, but in another, no atom is fully responsible for a single electron.

The formation of molecular clouds within galaxies is the first moment that molecules larger than dihydrogen appeared. Their binding energy is of course low by the standards of the warm interstellar medium, so only cold molecular clouds could be conducive to chemical reactions (because the frequency and mean energy of collisions is sufficient), while having temperatures low enough compared to binding energies that bonds would’nt immediately dissociate. Besides, highly energetic rays from other parts of the galaxies dissociates molecules, but the cloud’s opacity provides a degree of protection that alleviates that effect.

The formation of molecules in the gas phase can theoretically happen with direct reactions like: \[\text{O} + 2\text{H} \longrightarrow \text{H\_2O}\] But that is a slow process, since only a small fraction of the collisions will lead to a chemical reaction. Instead, the most common reactions will be exchange reactions where an ionized molecule can "donate" a proton. It all starts with ionized dihydrogen, which constantly gets generated when cosmic rays knock down electrons from neutral molecules: \[\text{H\_2\^+} +\text{H\_2} \longrightarrow \text{H\_3\^+} +\text{H}\] \[\text{H\_3\^+} +\text{O} \longrightarrow \text{OH\^+} +\text{H\_2}\] \[\text{OH\^+} +\text{H\_2} \longrightarrow \text{H\_2O\^+} +\text{H}\] \[\text{H\_2O\^+} +\text{H\_2} \longrightarrow \text{H\_3O\^+} +\text{H}\] \[\text{H\_3O\^+} +\text{e\^-} \longrightarrow \text{H\_2O} +\text{H}\]

The last step is dissociative recombination, which is extremely efficient when anion gets saturated with hydrogen. Simple hydrides like water (\(\text{H\_2O}\)), ammonia (\(\text{NH\_3}\)) and methane (\(\text{CH\_4}\)) are created this way. Later on, heavier atoms can combine via exchange reactions like: \[\text{CH} +\text{O} \longrightarrow \text{CO} +\text{H}\] This produces compounds like carbon monoxide (\(\text{CO}\)), carbon monoxide, (\(\text{CO\_2}\) , methanol (\(\text{CH\_3OH}\)) , etc. including molecules with dozens of atoms.

A key role is also played by dust, grains a few nanometers wide which formed around stars in the asymptotic giant phase, when solar winds are blowing elements from the outer shells away, or right after a nova event. When the surrounding medium is oxygen-rich, they tend to be silicate minerals like pyroxenes (\(\text{Mg\_x Fe\_{1-x} SiO\_3}\)) or olivines (\(\text{Mg\_{2x} Fe\_{2-2x} SiO\_4}\)). When carbon is more abundant, the dust is mostly carbonaceous, with a mix of pure crystalline carbon (i.e. diamond or graphite), and of amorphous hydrocarbons. In much smaller quantities we could finds metallic oxides, carbides and sulfides, as well as pure metals themselves. Dust not only plays a key role in the way it influences the thermal and optical properties of the surrounding medium, it also provides the surface where the most efficient chemical reactions occur.

As a result, a wide range of molecules have been detected in a variety of environments. Their origin is always interstellar (in molecular clouds) or circumstellar (in the envelope of giant stars, but nova events and nebulae have a way scattering these molecular products into neutral atomic clouds or beyond, into intergalactic space. Covalent bonds need more restrictive conditions to form, but then are remarkably resilient and we can find molecules in voids as well as near the very surface some stars. They range from simple diatomic molecules to the \(\text{C\_{70}}\) fullerene. The most common by far is carbon monoxyde \(\text{CO}\), with a general abundance of \(10^{-4}\) that of \(\text{H\_2}\) (for no other molecule is that figure higher than \(10^{-7}\)).

The stable future

15 billion years later, the universe’s ordinary matter is still made up of the same basic components: up and down quarks, electrons and neutrinos. Indeed, on such a scale they are almost perfectly stable. The half life of quarks has a lower bound of \(10^{34}\) years (ten decillion!). In a cool universe, quarks are bound into nucleons. Since neutrons decay into protons on much shorter scales, quark decay will largely happen in the form of proton decay, for which in most GUTs the principal reaction is \(p \rightarrow e^+ + \pi^0 \rightarrow e^+ + 2 \gamma\).

Crucially, we also know that the universe is electrically neutral (since the electromagnetic force has infinite range, like gravitation, but is much more intense, gravitational effects couldn’t dominate the macroscopic scales as they do, if it wasn’t the case). This means that the positrons will in turn disintegrate with electrons. The ultimate matter content of the universe will then be limited to neutrinos and potentially an LSP (like the neutralino). There will be no mediator bosons since those fermions are chargeless. As today, there will be one freely travelling boson: the photon.

States of matter

These changing conditions have a direct impact on structures at the microsopic level. We have previously discussed intramolecular bonds that hold atoms together within a molecule or a crystal lattice, with a typical strength between \(100\) and \(1000 \text{kJ/mol}\). The thermodynamic conditions of the stellar envelope enable the condensation of atoms in small metal chunks, ionic crystals or covalent networks, all bound by intramolecular forces, which would not have sustained the hotter conditions of the warm interstellar medium, not to mention of a star’s interior. As we’ve seen, these particles of dust, after being dispersed, change the thermal and optical conditions enough to create molecular clouds: pockets of cold interstellar medium where the low temperature and medium pressure, together with the capacity of the dust particles to act as a basic solvent and catalyze chemical reactions, promote the continued formation of covalent bonds.

Now we shall see an environment arise where a new class of forces becomes effective: inter-molecular forces, which act between molecules. Those are much weaker than their intramolecular counterparts (typically \(0.5\) to \(50 \text{kJ/mol}\)): inter-molecular bounds could not survive environments as hot as the stellar envelope, and could not emerge in pressures as low as found in molecular clouds (for a lack of encounters between molecules). The moderate temperature and significant pressure found in the proto-planetary disk, and later on the planets’ surface allows them to thrive.

Remember that we have characterized intramolecular forces as derived electromagnetic forces. They act between electrically neutral bodies (atoms), because those are not just neutral "points" in the ether, but have complex electronic structures. We are now going up another degree. Molecules also have zero electric charges, and have seemingly solved their electronic issues by filling or emptying their shells. The reason new forces emerge lies again in the uneven internal structure of the affected objects, this time their electrical polarity.

Consider water (\(\text{H\_2O}\)): the oxygen atom is more electronegative than the two hydrogen atoms, which means that it pulls the shared electrons closer to itself. The result is a dipole that, despite being electrically neutral overall, has a slightly negative pole (around the oxygen) and a slightly positive one (around the hydrogens). Note that the physical characteristic of the individual atoms play a role (like electronegativity), but also the geometry. The water molecule is bent, forming a triangle; if it was linear like carbon dioxide (\(\text{CO\_2}\)), it would not be polar.

There are different types of intermolecular forces:

The first category is not inter-molecular per se and blurs the line with intramolecular forces. It covers the various ion-dipole interactions, between a charged object and a polar molecule. These are stronger than any of the three forces above, since they involve a full charge rather than just partial ones.
Dipole-dipole interactions (often called Van der Waals forces) occur between polar molecules. They result from the attraction between partial charges: the positive pole of a dipole and the negative pole of another. In a large collection of dipoles, this results in a tendency to align, which of course is constantly disrupted by thermal agitation.

A specific case must be singled out for its particular strength, that of hydrogen bonds. This interaction occurs between a hydrogen atom "at the end" of a molecule within which it is covalently bound, and an outer electron pair from an electronegative atome in another molecule. Because hydrogen atoms are small and can approach the other dipole very closely, the force is much stronger than other Van der Waals forces, and in fact quantum effects play a role beyond simple electrostatic attraction. We will see what a considerable role hydrogen bonds play in chemistry and in particular biochemistry.
Lastly, London dispersion forces are the weakest, but also the most universal inter-molecular forces: they exist between all molecules, polar or not. At any given moment, the electrons in a molecule are not perfectly symmetrically distributed there is a fleeting, random dipole. That momentary asymmetry generates brief dipole-dipole-type interactions. Individually negligible, they add up significantly in large molecules with many electrons. For instance, they are the dominant force holding long hydrocarbon chains together, explaining how plastics can be solid at room temperatures while small molecules of the same family, like methane are gaseous.

Although they are weaker, the inter-molecular forces have a much momentous effect from our perspective as macroscopic observers. For a cosmic cloud to be made of individual atoms, of molecules, or of microscopic dust does not "feel different". But the adjunction of inter-molecular bounds gives rise to the familiar states of matter: solid, liquid, gas, upon which all the upper levels of organization are built.

How can that be? The key insight is that the inter-molecular forces are essentially attractive, and act as a counterweight to the statistical phenomenon that is thermal agitation: the manifestation of kinetic energy at any non-zero temperature that sees particles vibrating, rotating and jostling in random patterns.

When thermal energy vastly exceeds the strength of intermolecular forces, molecules can fly around freely: they fill whatever volume is available to them, take no fixed shape, and interact only through brief, elastic collisions. That state is a gas, the default state of matter in hot or sparse enough environments.

When the temperature drops, or the pressure decreases enough that intermolecular attractions can hold molecules in close proximity but not lock them into place, we get a liquid. Molecules are in constant contact, sliding past one another, forming and breaking intermolecular bonds continuously, a state of affairs often described as "short-range order, long-range disorder.". A liquid has a definite volume but no fixed shape it adapts to its container. It also exhibits a unique property: viscosity, its resistance to flow, which also depends on the nature and strength of the forces at play. Liquids held by weak dispersion forces flow almost like a gas, while honey, rife with a dense network of hydrogen bonds between the sugars, is very viscous.

Cool the system, or increase the pressure, even more, and thermal energy can no longer overcome the attractive forces: molecules settle into fixed positions, forming a solid. Note that the solid state was already present, on a microscopic scale, in the form of dusts made of tiny metal bits, ionic crystals (like salts) or covalent networks (like diamond). The cool, dense conditions of a planetary system enable the formation of larger solids held by these intramolecular forces. But when individual molecules settle into a solid, they can be held solely by inter-molecular forces. This is the case for instance when water turns to ice. There is no transfer nor sharing of electrons, as happens with intramolecular bonding: the molecules settled into place strictly on account of their inter-molecular bonding. All solids can be either crystalline or amorphous, which can depend a lot on the exact thermodynamic path followed to achieve freezing.

The transitions between these states melting, boiling, freezing, condensation, sublimation and deposition are called phase transitions and occur at specific combinations of temperature and pressure, which are characteristic of each substance and directly reflect the strength of its intermolecular forces.

Intermolecular forces’ principal contribution has then been to open the liquid state as a crucial middle ground, ultimately giving a medium to chemistry.

Chemical reactions can occur in clouds and on solid dust grains, as we’ve seen, but liquids are uniquely fertile, because they strike a balance: molecules are dense enough for frequent contact, but mobile enough to constantly rearrange and meet. Furthermore, they can act as solvents dissolve other substances, dispersing their molecules or ions evenly throughout the medium. The probability of productive encounters is then hugely amplified. The process of dissolution is itself governed by intermolecular forces. Polar solvents tend to dissolve polar or ionic solutes very efficiently, as all those dipoles start to interact. Nonpolar solutes like oil are better dissolved in nonpolar solvents.

Not all liquids were born equal. It is worth focusing on a very special one and examine some of its properties: I’m talking, of course, about water.

First, water is an exceptional solvent, due to its strong polarity. Take the simple case of common salt dropped in a glass of water. Salt is a ionic compound characterized by a very strong bond between its sodium cation and chlorine anion. Water molecules swarm their solute and form a shell around each salt particule, with a specific orientation imposed by the dipole-ion interactions. The collective strength of those interaction overcomes that of the ion-ion bound, and the two ions are divorced from each other, escorted by their respective water shells to be dispersed.

Water’s thermal properties include a very high specific heat capacity, which means that it takes significant energy to change its temperature, and a high vaporization heat, meaning that it resists boiling for longer after reaching its boiling point. Both results from its extensive hydrogen bonding network, which acts as a stabilizing force against thermal agitation. The consequences are momentous: at common pressures, water can be liquid over an extremely wide temperature range, and large bodies of water can absorb considerable amounts of energy, acting as a buffer. On a planet, oceans can prevent wild temperature swings between days and night, or between seasons.

Water also has an anomalous density behavior whereby ice is less dense than liquid water. As a result, ice water floats, which on a planet can insulate the liquid below (thus free to act as a solvent or a milieu for life). With other liquids, the solid phase would sink to the bottom, progressively freezing the whole body from the ground up.

Another of water’s talents comes from its autoionization. A small fraction of water molecules spontaneously dissociate into hydronium (\(\text{H\_3O}^+\)) and hydroxide (\(\text{OH}^-\)) ions,

Water also autoionizes: a small fraction of water molecules spontaneously dissociate into hydronium (\(\text{H\_3O}^+\)) and hydroxide (\(\text{OH}^-\)) ions, providing a baseline of reactive ions and enabling acid-base chemistry.

Last but not least: the hydrophobic effect. Nonpolar molecules hydrocarbons, oils, and the like do not interact favorably with water. Water molecules rearrange into a rigid, ordered cage around them which pushes the nonpolar molecules together, reducing the total surface area in contact with water. This is why oil drops coalesce in water: the hydrophobic tails huddle together, away from the water, while the hydrophilic heads face outward. This is the fundamental driving force behind the later formation of biological membranes from lipids.

Macromolecules

These properties, taken together, make water not just a good solvent but a uniquely hospitable medium for complex, multi-step chemistry, and later, biochemistry.

The Hadean environment can seem far from hospitable from our perspective: what qualifies as cool and loose compared to the surrounding of stars or warm cosmic clouds looks scorching hot and unbearably dense from our human standpoint. Add to that the frequent volcanic eruptions, constant outgassing, intense electric storms (millions of lightning strikes per day!), not to mention the rain of meteorites impacting the earth. Yet it may have been a kind of chemical paradise, all things considered. Although complex structure like living cells are much too fragile to sustain these conditions, the association of a liquid water environment (the primitive oceans), a chemically diverse mix, warm conditions and regular lightning, may have led to the appearance of the first macromolecules.

Macromolecules are large molecules, normally based on the repetition of at least one basic unit. Unlike structures like crystals or network covalent bonds, which are also based on the repetition of chemically-bonded structures, macromolecules are discrete entities with a clear beginning and end. Their physical characteristics and reactive properties can then be described in a predictable way. A "piece" of diamond is unique and not unlike any other, while two identical macromolecules are exactly interchangeable.

There is an almost complete overlap between naturally occuring macromolecules and organic molecules, that is, molecules whose mass is primarily from carbon atoms. Carbon is such a central element because of some key properties. Its tetravalency makes it a versatile atom from a bonding perspective. It easily bonds with hydrogen, oxygen, nitrogen,... all the most frequently encountered atoms in our planetary environment. It can form single, double or triple bonds, all of which have fairly high binding energy, with a diversity of geometrical configurations.

All macromolecules include strong, stable bonds between carbon atoms, which form chains, branches and rings. The simplest family of macrmolecules is hydrocarbons, which only includes \(\text{C-C}\) and \(\text{C-H}\) bonds.

Aliphatic hydrocarbons include:
- Alkanes: saturated chains with single \(\text{C-C}\) bonds only.
- Alkenes, with at least one double bond.
- Alkyles, with at least one triple bond.
Aromatic hydrocarbons, for instance benzenes, which exhibit a ring structures.

Other organic macromolecules can be thought of as derivatives of hydrocarbons, with at least one functional group tacked on: these are groups of atoms that determine the chemical and physical properties of the molecule. The main ones are (non-exhaustively):

Alcohol (hydroxyl \(\text{-OH}\) functional group bonded to a carbon atom)
Ether (\(\text{-O-}\) functional group bonded to two carbon atoms)
Aldehyde (carbonyl \(\text{C=0}\) functional group at the end of a carbon chain.)
Ketone (carbonyl \(\text{C=O}\) functional group in the midst of a carbon chain)
Carboxylic acid (carboxyl \(\text{O=C-OH}\) functional group)
Ester (\(\text{-COO-}\) functional group)
Amine (ammonia functional group where each hydrogen bond can be replaced with a full organic molecule)

As mentioned, a key consequence of macromolecule formation is the sheer diversity and complexity that immediately emerges. Bring together three atoms under the formula \(\text{CO\_2}\) and you get exactly one molecule, carbon dioxide. In terms of crystals or covalent networks, you get more diversity: we have seen that for identical chemical compositions, we can get allotropes: for instance, graphite and diamond are the two allotropes of pure carbon. But the constraints of unit cell repetition is a strong limit on diversity. Macromolecules, however, have many isomers, starting with relatively limited number of atoms. Two molecules are isomer to each other if they share the same atomic composition, but have different structural arrangements, leading to different physico-chemical characteristics. There are different types of isomerism:

Structural isomers have the same molecular formula, but different connectivity of atoms, for example different branchings in the hydrocarbon chains, different positions for the functional groups, or even different functional groups. For instance, \(\text{C\_3H\_8}\) has only one isomer, propane, but \(\text{C\_{18}H\_{38}}\) has 60,253.
Stereoisomers have the same connectivity, but different spatial arrangements.
- Geometric isomers differ in the angular arrangement of the 3D molecules, or in the order of specific connections around a given carbon atom (i.e. are two given connections contiguous, or on opposite poles).
- Optical isomers are non-superimposable mirror images of each other: they are called chiral molecules.

All the basic molecules from these families (methane, ethylene, acethylene, methanol, acetic acid, formaldehyde, cyanic acide...) have been observed in the galaxy’s molecular cloud. Some combination with several groups may have formed there too, but only when the chemically propitious medium of the early earth, and surely scores of other rocky planets, arose, could the infinite complexity of organic chemistry fully be expressed.

An important sub-class of organic molecules is that of biomolecules. Unlike organic molecules which, despite their name, are not necessarily related to life at all, biomolecules are the future building blocks of living beings. For the past three billion years, countless biomolecules have been synthetized by living cells. While it was established much earlier that biological matter was made of the same cloth, elements-wise, as inert stuff, it wasn’t at all clear that it could be created in the same way that inorganic compouds can be conjured until the twentieth century. That is when the famed Urey-Miller experiment, followed by others, showed that under conditions that roughly replicate that of the early earth (temperature, pressure, oceanic and atmospheric conditions, electrical sparks...), the building blocks of biomolecules can spontaneously appear.

The main classes of biomolecules are:

Amino acids are organic compounds that comprise an amine group and a carboxylic acid group. Long chains of amino acids are called proteins. An amino acid like glycin can be synthetized from formaldehyde, hydrogen cyanide and ammonia, with this kind of reaction:
\[\text{CH\_2O} + \text{HCN} + \text{NH\_3} \longrightarrow \text{NH\_2-CH\_2-CN} \quad \text{(aminoacetonitrile)} + \text{H\_2O}\] \[\text{NH\_2-CH\_2-CN} + 2\text{H\_2O} \longrightarrow \text{NH\_3} + \text{NH\_2-CH\_2-COOH} \quad \text{(glycine)}\]
Saccharides are a type of carbohydrate biomolecule. The simplest carbohydrate is glycolaldehyde (\(\text{HOCH\_2-CHO}\)), an organic molecule that contains both an aldehyde group and a hydroxyl group.It is readily derived from formaldehyde. Then, formaldehyde and glycoaldehyde together can react to form glucose, fructose and other monosaccharides. Those can then form macromolecular chains, forming disaccharides, oligosaccharides and polysaccharides, collectively know as sugars or starches.
Fatty acids are organic molecules with a carboxylic acid functional group attached to an aliphatic hydrocarbon chain, with the simplest ones being formic acid (one carbon, \(\text{HCOOH}\)) and acetic acid (two carbons). Long fatty acid chains form lipids, in general when several of them acquire a head, e.g. an alcohol like glycerol, a phosphate group or a sterol, and associate via esterification.
Nucleobases are nitrogenous organic molecules. They include purines (adenine and guanine) and pyrimidines (cytosine, thymine and uracil). For instance, adenin can synthetize from formic acid and cyanamide (\(\text{CN\_2H\_2}\)). Nucleobases combines with sugars (generally ribose or deoxyribose) to form nucleosides. In turn, nucleosides together with one or more phosphate groups can form nucleotides. One important example is adenosine triphosphate (ATP). Finally, nucleotides can form long chains, linked together through phosphodiester bonds. The resulting biopolymers are called nucleic acids, with the primary examples being of course deoxyribonucleic acid (DNA) and ribonucleic acid (RNA).

Biomolecules only exist in interstellar clouds in their simplest forms, like acetic acid, acetone, and perhaps glycin. The basic bio-blocks, quickly followed by their macromolecules versions likely formed on earth during the Hadean epoch.

The stable future

Biopolymers

We have seen how the building blocks of life amino acids, monosaccharides, fatty acids, nucleotides can spontaneously form from simple precursors, a process known as abiogenesis. Sensibly, the assembly of these monomers into long, functional chains gives rise to biopolymers. The fundamental reaction is conceptually simple: two monomers join together by releasing a molecule of water.

An amino acid links to another amino acid via a peptide bond.
A nucleotide links to another nucleotide via a phosphodiester bond.
Monosaccharides join through glycosidic bonds.
Fatty acids attach to glycerol backbones through ester bonds.

In each case, a water molecule is released. However these is an immediate thermodynamic issue with this. Under standard conditions, the formation of such bonds require energy. The opposite reaction, hydrolisis, which consists in water breaking the bond is favored.

So water is required as a solvent and a general facilitator for these chemical reactions to occur, but it is bound to actively resist, even revert, the polymerization! In proper biochemistry (involving actual living beings as an environment, energy carriers like ATP provide the necessary energy through their own hydrolisis. But in a primordial context, it is currently unexplained how widespread polymerization could happen. The main hypotheses center around certain clays or alkaline vents that may have created non-standard conditions where such reactions become favorable.

Proteins

A chain of amino acids linked by peptide bonds is called a polypeptide. A functional polypeptide, or an assembly of several polypeptides working together, is called a protein. The distinction is important: a polypeptide is defined by its sequence, while a protein is defined by its function, which depends critically on its three-dimensional shape.

Protein structure is usually described at four levels:

The primary structure is simply the linear sequence of amino acids.
The secondary structure refers to local folding patterns (shapes, stabilized by hydrogen bonds between backbone atoms from the main chain, and branches).
The tertiary structure is the overall three-dimensional shape of a single polypeptide chain, determined by the various interactions between the side chains.
The quaternary structure applies to proteins composed of multiple polypeptide subunits.

With 20 amino acids in use on Earth, and proteins having an average of 400 amino acids, the number of possible primary structures is already staggering. The possible higher structures make the potential diversity tremendously broader, but those are also not random. Proteins fold in a certain way in a matter of miliseconds as they are gently guided by weak forces between their components, first locally, then at higher and higher scales.

Nucleic acids

We have seen that nucleic acid is a long chain of nucleotides linked by phosphodiester bonds, and that each nucleotide consists of three parts: a sugar, a phosphate group, and a nucleobase. The sugar-phosphate units form the backbone of the chain; the bases project sideways, like the rungs of a ladder.

DNA (deoxyribonucleic acid) uses four bases: adenine (A), thymine (T), guanine (G) and cytosine (C), and normally exists as a double helix: two complementary strands wound around each other, with bases on opposite strands paired by hydrogen bonds A with T (two hydrogen bonds), G with C (three hydrogen bonds). This base-pairing rule is the foundation of genetic information storage: the sequence of bases along one strand completely determines the sequence along the other.

RNA (ribonucleic acid) uses uracil (U) instead of thymine and is usually single-stranded, though it frequently folds back on itself to form complex secondary and tertiary structures through internal base-pairing. RNA is chemically less stable than DNA (the extra hydroxyl group on ribose makes it more susceptible to hydrolysis).

Lipid membranes

Lipids are built from fatty acid chains attached to a head group. The crucial property of many lipids is that they are amphiphilic: one end (the head) is hydrophilic, the other (the fatty acid tails) is hydrophobic. Place them in water, and the hydrophobic tails cluster together, away from the water; the hydrophilic heads face outward, toward it. The most stable arrangement in most conditions is a bilayer: a dual sheet of lipids, naturally closing in on itself, forming a hollow sphere called a vesicle or, in biological context, a membrane.

This self-assembly requires no catalysts, no template, and no external energy: it is driven entirely by the hydrophobic effect and the thermodynamic tendency to minimize unfavorable contacts between water and the nonpolar tails. Laboratory experiments have shown that simple fatty acids can spontaneously form vesicles when agitated in water.

Polysaccharides

Polysaccharides are chains of monosaccharides with significant structural and chemical features. Cellulose, a polymer of glucose, is the most abundant biomolecule on earth and the structural backbone of plant cell walls, its rigidity arising from extensive hydrogen bonding between parallel chains. Chitin plays a similar role for arthropod skeletons. Starch and glycogen serve as energy storage molecules in plants and animals respectively.