S1C2: Chemistry

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a) Nature and relevance of chemistry

In Chapter 1 on Particles and Forces we took a quick tour of the inside of the atoms and the four fundamental interactions: electromagnetic, gravity, weak nuclear and strong nuclear forces. If you haven’t already read it, my advice is to do so, preferably now, since many indispensable concepts for the understanding of this chapter are introduced there.

Progressing up from the atom along the complexity chain, we are now stepping into the world of molecules and chemical compounds. The Encyclopedia Britannica definition for “Chemistry” is “the science that deals with the properties, composition, and structure of substances (defined as elements and compounds), the transformations they undergo, and the energy that is released or absorbed during these processes”. Understanding the atom, subatomic particles and fundamental interactions is key to the study of chemistry but once this knowledge is in place, the science of chemistry itself can then be used almost exclusively to understand a wide array of properties and behaviours which in turn can explain biochemistry and life, geology, pharmacology, and the manufacturing of various materials used in technical applications such as sulfuric acid, aluminium, LCD, Teflon and semiconductors.

The versatility of chemical compounds and processes is a consequence of emergence, a concept introduced in Section 1.b, with a handful of variables resulting in a staggering multiplicity of properties and optionality when it comes to creating, transforming and interacting with more complex structures. The main variables are the position and charge of the electron, the composition of the nucleus, and the internal energy of a system. By the end of this chapter, I invite you to reflect back on how many of those properties and behaviours you think could easily have been extrapolated and how many are downright “unexpected”, with most of them somewhere in between. Evidently, it is much easier to work backward and explain the reasons and steps behind a specific phenomenon than to predict one that is as yet unobserved but could in reality genuinely eventuate, in the sense that the predicted structure or system is stable enough.

This explosion of options and possible behaviours will take on even more incredible proportion as we move into the realms of biochemistry and then biology. And so, understanding chemistry is the second step in this journey.

b) Molecules, ions and chemical bonding

When two or more atoms come together to create a stable structure, they form a molecule, which may or may not consist of identical chemical elements. For instance an oxygen molecule (O₂) consists of two oxygen atoms and carbon dioxide (CO₂) is made of one atom of carbon bound to two oxygen atoms. Those are fairly simple molecules though and they can grow to become quite complex, especially in the case of organic compounds. Take the example of benzene (C6H6): it has 12 atoms in total, 6 carbon (each bonded to 2 others to form a ring) and 6 hydrogen (each bonded to one carbon atom). So, how do these molecules come to be?

An atom or a molecule with a net negative charge is called an ion. If it is positive (a “cation”), this means it is missing some electrons because those are negatively charged. Conversely, an “anion” has extra electron(s) and is thus negatively charged; in other words the charge comes from the imbalance between the number of negatively-charged electrons and positively-charged protons. Unsurprisingly, ions with opposite electric charges are pulled towards each other due to electrostatic force, an instantiation of the electromagnetic interaction, and can form ionic compounds held together by “ionic bonds”. The other name for an ionic compound is salt, so the most obvious example of this category of molecule is perhaps sodium chloride (NaCl), better known as edible salt.

However, ionic bonds are not the most prevalent form of naturally occurring chemical bonds, this distinction belongs to covalent bonds. The name originates from the concept of valence electrons, those particles located into the outermost shell of an atom – in the days of the graphical representations based on the Bohr’s atom, this would have been the orbit furthest from the nucleus. The notion of atom shells pertains to the domain of quantum mechanics, so for purpose of this chapter and chemistry in general we will retain the following idea: there is a maximum number of electrons a specific shell can accommodate and, if it is not filled up, these can be involved in the formation of a chemical bond. In covalent bonds, two atoms share a pair of electrons thus there is no net effect on their respective electrical charge; examples include the Nitrogen molecule (N₂) or Water (H₂O) where the oxygen atom has two covalent bonds, one with each hydrogen atom.

There are a few exceptions such as one- and three-electron bonds (not covered in this chapter, nonetheless I have included a hyperlink to the covalent bond entry on Wikipedia in the last section of this chapter if you wish to learn more) and a variation called delocalized bonding, which is the sharing of electrons over more than two atoms. Delocalized bonding is the most frequent basis for the constitution of metals and some of their best-known properties such as high thermal and electrical conductivity. The term used in this case is metallic bonding and if you want to represent it in your mind’s eye it would be the sharing of a cloud of free electrons by metal cations (positive ions of elements classified as metals).

Finally, the fourth main type of chemical bonding is called hydrogen bonding, it consists in an electrostatic attraction between a hydrogen atom and another atom called “acceptor atom” from another atom with a “free pair” of electrons. A free pair refers to two electrons located on the last shelf that are not already involved in a covalent bond; atoms with such profile tend to be more electronegative, meaning they attract shared electrons. However, since there is no sharing of electrons involved per se, this type of bond is much weaker than a chemical bond. Note that there are a few more types of chemical bonding but they are less frequent and central to the grand edifice leading up to life so I will leave it at that as far as chemical bonds go.

Figure 2: Graphical representations of Atisane molecule

Credit: uploaded by Ddoherty on Wikipedia (CC BY-SA 3.0)

The above 3D (left and centre) and 2D (right) representations have been included to convey two ideas: the first is that molecules have defined shapes in space, something that is important when it comes to interactions with other compounds, whether it is their propensity to react or create larger compounds, and the second is the need to have a formal scalable way to specify the structure of a molecule so that any of them can be abstracted in a consistent fashion and some of its properties be made readily apparent. The 2D diagram on the right is called a skeletal structural formula, as opposed to the molecular formula C20H34, it exhibits the spatial organization from which it is mostly possible to deduce the type of chemical bonds between each elements.

Other types of chemical formulae can be used depending on the type of information one wishes to convey. For example, the molecular formula for butane is C₄H₁₀, the condensed formula for the same molecule is CH₃(CH₂)₂CH₃, the empirical formula of C2H5 indicates the ratio between elements, and the diagrammatic structural formula can be described as follows: 4 carbon lined up horizontally, the first C atom on the left is also linked to hydrogen atoms at the top, left and bottom and by symmetry the last carbon atom on the right has hydrogen atoms at the top, right and bottom, whereas the two central C atoms each have links to two hydrogen atoms above and below them. When it comes to ions, the positively charged cation is noted with a “+” exponent as in Na⁺ for Sodium and the negatively charged anion as in SO₄²^– for sulphate, which has 2 extra electrons.

c) Chemical elements

This feels like the right time to properly introduce the concept of chemical element, a notion already used several times in this chapter and the previous one. It can be defined as a chemical compound whose individual atoms all have the same atomic number, a term designating the number of protons in the nucleus of an atom. When the number of neutrons changes, we refer to isotopes of this element.

Protons have been selected over neutrons as the referenced particle for atomic number not on the basis of either their size or weight but because of their electric charge, which has implications in terms of electron shells and electron valence, and therefore in terms of properties. So much so that the observation of a recurring pattern of similar properties across different elements led to the idea of elements periodicity, with the number of protons as the key variable. In search of an organizing principle, scientists eventually worked out the periodic table of the elements, a simple example of which is enclosed below though there are more comprehensive colour-coded versions arranged by blocks such as transition metals for the d-block and the diagonal dividing line between metals and non-metals. Those blocks refer to atomic orbitals, a concept closely related to the electron shells, and the numbered columns refer to groups (all elements outside of the f-block) comprising elements with similar chemical or physical characteristics. The rows are called “periods” and the number in front of them designates the number of electron shells.

So powerful was this arrangement that it had predictive powers in terms of where to find new elements and their atomic weight. Dimitri Mendeleev was the chemist who formulated the Periodic Law and created the first accepted version of the periodic table of the elements, in 1869.

On Earth, elements up to 94 plutonium can be found whereas the rest (up to number 118 Oganesson) were synthesized within laboratories though some were also observed in distant stars through spectroscopy. Of note, all transuranic elements (beyond 92 uranium) are radioactively unstable and will over time lose energy by radiation, eventually decaying into other elements through nuclear transmutation. It is no coincidence that enriched isotopes of uranium (U-235) and plutonium (Pu-239), some of the heaviest naturally occurring elements, were selected as fissile materials for nuclear weapons.

d) Intermolecular bonding and macromolecules

We now have a decent grasp of the nature of chemical elements, of their propensity to bond with and react to other elements, and of how atoms/ions link with each other to form molecules. Hence, we are sufficiently equipped to look into how some of those same bonding mechanisms and other type of interactions between molecules give rise to larger, more complex entities.

Enters the macromolecule, the bedrock of biochemistry – the central topic of upcoming Chapter 4. The term of macromolecule is quite loosely defined and really refers to a very large molecule one could picture as an agglomeration of simpler molecular units held together by covalent bonds, i.e. the sharing of electrons. There is no limit to how large such an assemblage can be and some macromolecules comprise several thousand atoms.

The most common type among them is called polymer and consists of repeating molecular units called monomers; those can be natural as in the case of proteins and latex rubber, or synthetic. Examples of the latter are ubiquitous in our daily lives in the form of plastics such as polyethylene (think plastic bags) or polyvinyl chloride (PVC – used for pipes). It is their nature as polymer that endows these materials with some of their most salient properties such as toughness (the ability to absorb energy) and elasticity (defined by Britannica as the ability of a deformed material body to return to its original shape and size when the forces causing the deformation are removed).

The next steps in this game of Lego devised by nature is the formation of structures involving several molecules. If my explanations so far were sufficiently clear, it should be obvious that covalent bonding is not involved here (otherwise we would be dealing with macromolecules) and therefore that intermolecular forces are weaker than the intramolecular ones. We already learned about hydrogen bonding and although I did not point it out then, it is a form of intermolecular bonding since the electrostatic force responsible for this type of link takes place between a hydrogen atom and another acceptor atom belonging to a different molecule. There exist several other types of intermolecular forces like van der Waals forces or dipole-dipole but these get quite technical, and the main ideas to take away are the very existence of these interactions between molecules and the emergence of different types of properties out of these. Hence, I will skip an in-depth analysis of intermolecular forces and instead make a couple of generic observations we can draw from the previous sections.

Firstly, all the intra-molecular and inter-molecular interactions involve the electron – talk about a pivotal role in the universe. This is a function of two variables: the respective strength of fundamental forces at the atomic level and the position in space of the electrons, far from the nucleus and “orbiting” in electron shells forming the “outside edge” of the atom so that they represent the most accessible and potent interlocking device.

Secondly, the nature of intra-molecular interactions dictates the shape and structure of molecules and, together with inter-molecular interactions, it is involved in determining the states of matter.

e) States of matter and phase transitions

States of matter can loosely be defined as the way a specific type of matter is arranged and its constituting units (atoms/ions or molecules) can move or not around each other. The four states commonly found in nature are solid, liquid, gas and plasma. Let’s explore these in turn.

In solids, the forces between particles is strong enough to prevent them from moving so they have a stable shape and set volume. Accordingly, they exhibit structural rigidity and can resist forces applied to their surfaces. Solids can undergo phase changes to and from liquids or gases. The phase transition from solid to liquid is called melting and the reverse is called freezing. Although it is not the most commonly observed transition in our daily life, the change from solid to gas is called sublimation and the other way around is deposition.

Compared to solids, the particles in a liquid are not fixed; they are still held together by inter-molecular forces but the bonds between specific particles are only temporary and keep shifting. This is the reason behind the fluidity of a liquid, as compared to the rigidity of a solid. The phase transition from liquid to gas is called evaporation (or vaporization) and the opposite is termed condensation.

In gases however, the kinetic energy of particles (related to their motion) dominates over inter-molecular forces, hence the particles are not bonded together, unlike a liquid or solid. This means gases are not only fluid but also compressible or expandable.

The fourth state of matter we can readily observe is plasma. It is made up of a high concentration of charged particles, both ions and electrons, thus it is particularly electrically conductive. The phase transition from gas to plasma is called ionization and the reverse is recombination. We can witness plasma in lightning, fluorescent light bulb, in some science museum exhibits, and more importantly in stars and clusters of matter within and between galaxies. In the scheme of the entire universe, these outer space volumes completely dwarf the solid, liquid and gas states: plasma accounts for about 99% of ordinary matter. Not what our intuition tells us clearly, an obvious example of sampling bias.

There are a few more esoteric states of matter such as liquid crystals that mix properties of ordered solids and some particles fluidity, glass which is technically a solid lacking the long-range order of crystals, superconductors with no or near-zero electrical resistivity, and several more in high-energy states. For those who are curious, I have enclosed a link in the last section to the Wikipedia entry with a comprehensive list of the states of matter.

To conclude this section, it is worth diving into the why of phase transitions, their trigger. The answer is already partly available when we recall the interplay between kinetic energy, intermolecular forces and the other types of bonds involved in the “fixity” of particles in the solid state. The first driver of phase transition is a shift in pressure: the higher the pressure, the shorter the distance between particles and the stronger the bonding forces experienced by particles. At high pressure, a liquid can be turned into a solid and, of course, a gas into a liquid. This explains why, all else being equal, your saucepan full of water boils at 100° C around sea level (exactly at 1 atm or 101,325 pascals) and as you camp at high altitude the boiling point is lower so that your tea doesn’t infuse as well and the water purification isn’t as thorough. The second key variable is temperature: the higher it is the more kinetic energy in the system and the less effective the intermolecular forces are in holding particles together. Heat up a solid and it becomes fluid, either a liquid or a gas, depending on the pressure.

f) Chemical reactions

Thus far, we have examined the creation of molecules and macromolecules, the different types of chemical bonds linking atoms, ions and molecules, how variables like pressure and temperature influence distance between particles and the interplay between kinetic energy and intermolecular forces impacting the arrangement of those entities and the fluidity of matter. Nevertheless, we are still to look into what happens when those chemical bonds break? And why does this happen?

The simple answer to the first question is similar to the unmaking of Lego blocks, this renders the blocks available again to build another structure. Likewise, the breaking of chemical bonds make the electrons available again to form new bonds, which under some circumstances will lead to a rearrangement of the atoms. In other words, there will be changes in the type of chemical substances present before and after this event called a chemical reaction. The obvious key requirement is that there should be at least two types of chemical substances involved (except in the case of radioactivity where there will be changes to the nucleus itself so one element breaks down into other elements); the original substances are called reactants or reagents, and the output of the reaction will be one or several products (this is the jargon term). It should be noted that most of these reactions are reversible.

If we are to sum the energy locked or captured within the various inter-particle bonds before and after a reaction, we find energy levels will have changed. However, since the total energy in a system is constant (this is the law of conservation of energy we will discuss in Section 6.d), this implies energy is either released by or required for the reaction, typically in the form of heat. An exothermic reaction is one that releases heat – combustion for example, the converse is called an endothermic reaction and it necessitates energy to be brought into the system to be triggered. This explains why reaction rates are positively correlated with temperature.

It is often the case that chemical reactions involve an increase in the oxidation state (loss of electrons) of one element and a simultaneous reduction in the oxidation state of another element. An increase in oxidation corresponds to a loss of electrons and the technical term of “redox” refers to the opposite, a reduction in oxidation. Redox can occur via electron transfer or atom transfer. For background, oxidation often involves oxygen but not always, though Lavoisier thought it did when he worked on combustion (a type of redox reaction) in the 1770s.

Sometimes the products of a first reaction will enter into a second reaction yielding another set of products (different from the original ones, so not a reversed reaction). This requires the presence of at least 3 sets of chemical elements and it is the process upon which catalysis relies, a technique we frequently resort to in order to accelerate the rate of specific chemical reactions. By introducing a new chemical compound called catalyst, one of the original compounds reacts with it to yield “intermediates” that then react with another of the original compounds to yield the desire product and the catalyst back in its original form, so the latter can be reused. This speeding up occurs because the actual type of chemical reactions involved are different from the reaction that would have taken place without the catalyst.

I trust this and the previous sections showcase the versatility of creation and transformation underpinned by the fundamental forces, the laws of thermodynamics and our beloved electron.

g) Trivia – Noble gases

Noble gases, also called inert gases, occupy the last column of the periodic tables and are so named because of their low reactivity. At the end of the 19th century, it was indeed considered a quality of the well-bred not to lose one’s dignity when being provoked (at least not on the spot).

The reason for their low reactivity has to do with their outer electron shell being filled up or “closed” with 8 electrons, except for helium since the first shell only has two spots available; this configuration is not propitious to chemical bonding with other atoms or ions. This makes noble gas a safe choice as fillers for various applications such as helium in airships or argon in incandescent light bulbs. Their weak interatomic propensity to bond means they also have low boiling and melting points, a useful property for use as refrigerants. Finally, under standard conditions, we perceive them as lacking colour, odour and taste.

The six noble gases are helium, neon, argon, krypton, xenon and radon with oganesson also technically in that category but being both synthetic and unstable.