The Graphene Revolution

by Brian Clegg

  Not uncuttable

  By the start of the 20th century, just as atoms were coming to be accepted as real things, evidence was showing up that atoms really didn’t live up to their name as being ‘uncuttable’. It all started with the unlikely discovery of cathode rays. The term may raise vague memories of old TV sets and computer monitors – the ones with the big lump sticking out at the back. Found in most homes up to the 1990s, these made use of what was essentially the same technology as that used by the likes of Victorian British physicist William Crookes. Crookes was a self-taught scientist who did early work on electrical effects in vacuum tubes, sealed glass tubes with most of the air pumped out. These tubes were commonly referred to as Crookes tubes.

  Experimenters had been putting electrical charge across two electrodes inside such a partially evacuated tube since Michael Faraday noticed this produced a strange glow in the 1830s, but as better vacuum pumps became available, it was possible to remove most of the gas from the tube and the result was that the bulk of the tube went dark, but something invisible passed down the tube and caused the glass at the end to glow. In the classic Crookes tube demonstration, the flow is from a negatively charged cathode, down the tube, past a positively charged anode to hit the glass, with the anode’s shape (often a Maltese cross, for some reason) left dark as a shadow.

  As researchers became more familiar with these ‘cathode rays’, they painted the end of the tube with a substance such as zinc sulfide, which was more fluorescent than glass, producing a brighter glow. The CRT (cathode ray tube) TVs were simply more sophisticated versions of this, where the ‘cathode rays’ were steered by magnets and electrical fields to create a picture. But what were these rays? Some thought they were charged matter, atoms that had picked up an electrical charge (what we’d now call ions), while others thought they were a different form of electromagnetic radiation.

  The Cambridge-based physicist J.J. Thomson managed to measure the mass of the particles that made up this ray, showing that they certainly weren’t a form of light, using a combination of the heat they produced when hitting a metal junction and the amount they were deflected by magnetism. However, his result suggested they weren’t ions either. Thomson found that the ray’s components were at least 1,000 times lighter than the smallest atom, hydrogen, and later refined this to around 1,800 times lighter than hydrogen.

  These particles, which seemed to be coming out of stray atoms of gas or the matter of the electrodes – and so were emerging out of atoms – were much smaller than the atoms themselves. It appeared that the individual atoms were being cut – or at least that tiny parts of them were being pulled off. These particles, which Thomson called corpuscles, were soon better known as electrons, a name that had already been used as the basic unit of electrical charge produced by a battery.

  Of itself, the existence of electrons did not say too much about what was going on inside an atom. Just because something comes out of a black box, that doesn’t mean we know what’s going on inside. Thomson’s own theory, often called the plum pudding model, was that atoms were made up of a collection of negatively charged electrons, scattered through a massless, positively charged ‘matrix’ that held them in place electrostatically, making the electrons the ‘plums’ ¶ in the model.

  One aspect of Thomson’s model that now seems strange is the idea that this positive matrix had no mass. This meant that if Thomson’s model was accurate, all the mass of the atom had to come from its component electrons. With modern figures for the mass of atoms, this would have meant that hydrogen, an atom we now know contains a single electron, would have required a total of 1,837 electrons in in order for it to have sufficient mass.

  A more useful picture of the atom would emerge from Manchester, where the New Zealand-born physicist Ernest Rutherford had his lab. Rutherford’s team were experimenting on alpha particles – positively charged particles with a mass closer to that of an atom than that of an electron, and which were emitted by newly discovered radioactive materials. Before he had arrived in Manchester, Rutherford had discovered two different types of rays, which he named alpha and beta rays, later renamed as particles. Each was electrically charged, but with opposite values, curving in different directions if passed through electrical fields. Alpha particles would later be identified as the nuclei of helium atoms, each particle consisting of two protons and two neutrons (the particles that make up the atomic nucleus), but this was not known at the time. Similarly, beta particles turned out to be high-energy electrons.

  After a number of preliminary experiments, in 1913 Rutherford’s team had set up an experiment whereby a stream of alpha particles from a radioactive source (radon gas) was directed towards a thin piece of gold foil. The experiment took place in a cylinder from which the air had been removed. An observer – usually either Hans Geiger || or Ernest Marsden, Rutherford’s assistants, had to sit in the dark in low light conditions until their eyes had acclimatised and then peer through a microscope towards the gold foil. The end of the microscope had a zinc sulfide panel fixed to it, so any alpha particles heading in that direction would cause a tiny flash. The observer would watch at different angles, rotating the microscope between observations to detect how many, or indeed if any, particles were deflected.

  The Manchester physicists expected that when the alpha particles came close to the atoms in the gold, some would be slightly deflected by interacting with the electrical charges in the atom – which was the case. But a totally unexpected result was that some of the particles bounced back near to the direction they came from. In one of his more famous quotes, Rutherford commented that ‘it was as if you fired a fifteen-inch shell at a piece of tissue paper and it came back and hit you’. This could only have occurred if the positive charge in the gold atoms, rather than being the diffuse matrix that Thomson imagined, was concentrated into a small, dense core containing much of the atom’s mass. **

  It says something for Rutherford’s team that they took the precaution of looking for alpha particles travelling in a totally unexpected direction. A less capable experimenter might well have failed to do this and would have missed the breakthrough observation. In any case, Rutherford stole a piece of terminology from biology and called this the nucleus of the atom.

  The solar system model

  The Manchester discovery led to the beguiling idea that atoms were similar to miniature versions of the solar system with the nucleus at the centre of the atom taking the role of the Sun and the electrons, flying around the outside, playing the parts of the planets. There was a pleasing sense of symmetry if this were the case – and there’s nothing physicists like more than symmetry. However, there was a fundamental problem with this model of the atom, which was never taken seriously by the physics community, despite still tending to be used to this day as the graphic art representation of an atom.

  There’s one big difference between a solar system and an atom. The force between the central massive body of the star and the orbiting planets in the solar system is gravity – but in an atom, the force between the nucleus and the electrons is electromagnetic. And these two forces of attraction don’t work the same way. In simple terms, you can’t keep something in orbit using electrical charge.

  It was already known by Rutherford’s time from the work of the great Victorian Scottish physicist James Clerk Maxwell that whenever you accelerate an electrical charge, it gives off energy in the form of electromagnetic radiation. This is how radio works, for instance. The signal from the transmitter accelerates electrons up and down the aerial, losing some of the electrons’ energy in the form of photons – electromagnetic radiation. However, to stay in orbit around the atom’s nucleus, electrons would be constantly accelerating.

  It doesn’t seem that this is the case, as an orbiting body usually moves at a constant speed. However, there is still acceleration, because this is any change of velocity , which is a measure of both speed and direction of travel. To keep in orbit, the direction of movement is always chan
ging. That’s fine with gravity in charge, but it would mean that the orbiting electrons would quickly lose their energy in a flare of electromagnetic radiation and plunge into the nucleus. Every atom would collapse.

  To solve the mystery of the atom, other physicists employed the new and rapidly developing quantum theory – of which more later, as it will be essential to explore some of the stranger properties of graphene. The quantum model of the atom began with the young Dane Niels Bohr, who, rather neatly, started work on the problem while spending a year working with Rutherford in Manchester. All we need for the moment, though, in our understanding of how atoms interact to form material structures, is the idea that atoms have a small, dense, positively charged nucleus and are surrounded in some way by one or more moving electrons, plus the additional observation of the relationship between atomic structure and the chemists’ periodic table.

  The idea of putting the different elements into a structured table had been around for some time before the Russian Dimitri Mendeleev produced his earliest attempt in 1869 at what we now know as the periodic table, grouping elements into columns based on similar behaviour, with increasing atomic weight as you move down the table. As the structure of atoms became better known, it seemed clear that electrons occupied one or more ‘shells’ †† around the atom, each of which could only have a certain number of electrons in it. How they occupied it without atomic collapse was not yet clear.

  The chemical behaviour of an element, it was discovered, crucially depended on the number of electrons, and the number of open spaces, in the outermost occupied shell. You may remember something from school science called ‘valence’ which describes how an element is likely to combine with other elements to form compounds – this is a direct reflection of the state of that shell. So it is the detail of the atomic structure which gives chemists and physicists, whether dealing with multiple versions of the same element or combinations of different elements, a mechanism to describe how they will interact with each other.

  Bonding sessions

  Whether they were supporters of the four/five element theory or atomists, early natural philosophers had recognised that something must enable different elements or atoms to stick together so that they could make up the more complex stuff that we experience all around us (not to mention the complexities of our own bodies). Even the simple model of wood being made of earth, air, fire and water (see page 21 ) required this. By Newton’s time, there was speculation that these connections between the elements were physical links, perhaps due to the shapes of atoms (for those who believed in atoms), or the result of some type of inter-elemental glue. But Newton, with the success of his theory of gravity, and familiar with the effects of magnetism, preferred the idea that there was some kind of attractive force linking the component parts together.

  Inspired by Rutherford’s model of the atom and gradual realisation of the role of electron shells, the American chemist Gilbert Lewis had by 1916 come up with the idea of an electron pair bond, or covalent bond, linking atoms. This was where one or more electrons in an outer shell, instead of belonging to a single atom, were shared between two atoms, forming a bond between those atoms due to the electromagnetic attraction between the electron and the two atoms’ nuclei. Effectively, each atom made a claim on this electron or electrons as part of its structure. The electrons were Newton’s glue.

  The same year, German physicist Walther Kossel came up with a different way that atoms could bond where the outer shell had one more or one fewer electron than the normal atom. The result would be that the ‘ion’ ‡‡ with the extra electron would be negatively charged and the ion that was missing an electron would be positive – so the two could be electromagnetically attracted together in an ‘ionic’ bond. Once again, Newton had had the right idea with the concept of attraction.

  It’s the existence of these two types of bond that not only makes it possible for chemical compounds to exist – from the ionic bond forming the simple sodium chloride that is common salt to the vast array of covalent bonds in the huge molecules of DNA – but, even more fundamentally, allows nature to go beyond individual atoms or molecules to produce physical objects made up of many billions of atoms, all linked together by different kinds of bonds. It’s thanks to bonds that we are able to have solid materials. And as we come to look at graphene’s capabilities, the nature of its bonds will be an important part of the way it performs.

  Solids and structures

  It is something of a self-evident fact that not every solid is the same, even when made out of the same type, or types, of atom. The way the bonds form between atoms, and the structure that the arrangement of bonds produces, helps determine the substance’s physical properties – not just how it looks but how it reacts with other substances, its melting point, its strength and far more. As we shall see, it is the particularly impressive structure that carbon is capable of forming in graphite that results in graphene’s remarkable properties.

  Broadly speaking, solid substances tend to be either crystalline or amorphous. As we have seen, a crystalline substance is any one that has its atoms or molecules bonded together in a regular, recurring lattice. Many solids, from salt crystals to metals, do have such structures, but in others the bonds are higgledy-piggledy, without any repeating structure – these are amorphous solids, such as glass and many of the plastics.

  However, while some atoms or molecules always form solids with the same kinds of structure, others have a range of options available to them – few more so than carbon – and the shapes formed by the bonds linking the molecules can have a crucial effect on the substance’s physical characteristics, such as its strength, melting point and electrical conductivity.

  A familiar example of the impact of the lattice shape on the physical characteristics of a substance is provided by solid water – or ice, as it is better known. Each individual water molecule does not have its three component atoms in a straight line, but instead has a wide-angled V-shape, §§ with the hydrogen atoms at the top points of the V and the oxygen at the bottom. This shape, combined with the attraction between the relatively negative oxygen atom in one molecule and the relatively positive hydrogen in another – an attraction known as hydrogen bonding – makes it fairly easy for the water molecules to form crystals surrounding a hexagonal space (and is responsible for the six-pointed nature of snowflakes).

  Because of that particular shape and angle between the bonds, the lattice they form is not the most tightly packed that they can be. Water molecules can get closer together in a low-temperature liquid than they are when the solid crystal forms. This means that as water freezes it has the unusual (although not unique) property of becoming less dense as a solid than it was as a liquid. This means that ice floats on water (and tends to burst through containers that it is frozen in). ¶¶ The simple hexagonal form of normal ice, incidentally, is by no means the only structure that water can adopt as a solid. There are at least seventeen different structures it could produce, but at usual freezing temperature and under the Earth’s atmospheric pressure, the hexagonal form dominates.

  Carbon is also able to solidify in a range of structures, which enable exactly the same element arranged in different configurations (known as allotropes) to behave as if they were unrelated substances. The best-known allotropes of carbon are diamond, which has an interlocking cube-shaped lattice of atoms, giving it great strength, and graphite, which, as we’ve already discovered, is made up of atom-thin layers of repeating hexagonal structures, individual layers being known as graphene. It’s also possible to have carbon structures that make up relatively small molecules with closed structures, their bonds being like the pentagons and hexagons that make up the lines on the outside of a football.

  These closed forms are known as fullerenes (or buckeyballs as a less serious nickname), both references to the American architect Buckminster Fuller, who designed geodesic domes that were similar to parts of a fullerene. The best-known buckeyball molecule, buckminsterfullerene,
has 60 carbon atoms in its structure. Another, more open form of fullerine consists of a tube made of the same flat carbon lattice as graphene, wrapped around to form a cylinder. Such ‘carbon nanotubes’ are, in effect, tiny tubular pieces of graphene (think of taking a piece of paper and rolling it to produce a tube). The carbon fibres embedded in a polymer in everything from car dashboards to bicycle frames, producing a material that is misleadingly usually just called ‘carbon fibre’, may well contain some nanotubes, although they are mostly just strands of carbon chains, a bit like multiple thin strips of graphene. Although mostly artificial, fullerenes can occasionally occur in nature.

  Diamond

  Graphite

  Buckminsterfullerene

  Two further allotropes of carbon are significant. One is lonsdaleite (named after the British crystallographer Kathleen Lonsdale), which is like diamond but has a hexagonal, graphite-like lattice instead of the usual cubic lattice. Lonsdaleite was first found in a meteorite and has also been artificially produced by putting graphite under high pressure and temperature. It should, in principle, be even harder than diamond, though existing specimens have tended to have a lot of impurities and more incomplete lattices than a good-quality diamond, making it weaker than its more familiar cousin. The other, less fancy allotrope is amorphous carbon, which lacks a uniform lattice structure. This is most familiar as coal or in the flecks of carbon that make up soot.