50 Chemistry Ideas You Really Need to Know

Book 0: 50 Ideas You Really Need to Know series

by Hayley Birch

Genres: Short Stories

Synopsis

Master the chemistry ideas that shape the world we live in today.

In a series of 50 accessible essays, Hayley Birch introduces and explains everything you need to know about the world of chemistry, offering fascinating insights into our origins and life as we know it.

From the molecules that kick-started life itself to nanotechnology and from fermentation to the periodic table, 50 Chemistry Ideas You Really Need to Know is a complete introduction to the most important chemistry concepts in history.

Contents include: Thermodynamics, Catalysts, Fermentation, Green Chemistry, Separation, Crystallography, Microfabrication, Computational Chemistry, Chemistry Occurring in Nature, Manmade Solutions: Beer, Plastic, Artificial Muscles and Hydrogen Future.

Release date: January 1, 2015

Publisher: Quercus Publishing

Print pages: 208

Sample

Atoms are the building blocks of chemistry, and of our Universe. They make up the elements, the planets, the stars and you. Understanding atoms, what they’re made of and how they interact with each other, can explain almost everything that happens in chemical reactions in the lab, and in nature.

Bill Bryson famously wrote that each one of us might be carrying up to a billion atoms that once belonged to William Shakespeare. ‘Wow,’ you may well think, ‘That’s a lot of dead Shakespeare atoms.’ Well, it is and it isn’t. On the one hand, a billion is about the number of seconds that each of us will have lived by our 33rd birthday. On the other hand, a billion is the number of grains of salt that would fill an ordinary bath, and less than one billionth of one billionth of the number of atoms in your entire body. This goes someway to explaining how small an atom is – there are over a billion times a billion times a billion just in you – and suggests that you don’t even have enough dead Shakespeare atoms to make up one brain cell.

Atoms are so tiny that, until recently, it was impossible to see them. That fact has changed with the development of superhigh resolution microscopes, to the point where, in 2012, Australian scientists were able to take a photograph of the shadow cast by a single atom. But chemists didn’t always have to see them to understand that, at some fundamental level, atoms could explain most of what goes on in the lab, and in life. Much of chemistry is down to the activities of even tinier, subatomic particles called electrons, which make up the atom’s outer layers.

If you could hold an atom in your hand like a peach, the stone in the middle would be what is called the nucleus, containing the protons and neutrons, and the juicy flesh would be made up of electrons. In fact, if your peach was really like an atom, most of it would be flesh and its stone would be so small you could swallow it without noticing – that’s how much of the atom is taken up by electrons. But that core is what stops the atom drifting apart. It contains the protons, positively charged particles that hold just enough attraction for the negatively charged electrons to stop them flying off in all directions.

Atomic theory and chemical reactions

In 1803, the English chemist John Dalton gave a lecture in which he proposed a theory of matter based on indestructible particles, called atoms. He said, in essence, that different elements are made of different atoms, which can combine to make compounds, and that chemical reactions involve a rearrangement of these atoms.

Not all atoms are the same. Extending the fruit analogy further, atoms come in many different varieties, or flavours. If our peach was an atom of oxygen, then a plum might be, say, an atom of carbon. Both little balls of electrons surrounding a proton pip, but with completely different characteristics. Oxygen atoms float around in pairs (O2) while carbon sticks together en masse to make hard substances like diamond and pencil lead (C). What makes them different elements (see page 8) is their respective numbers of protons. Oxygen, with eight protons has two more than carbon. Really large, heavy elements like seaborgium and nobelium have more than one hundred protons in each of their atomic cores. When there are this many positive charges crammed into the vanishingly small space of the nucleus, each repelling the other, the equilibrium is easily upset and heavy elements are unstable as a result.

Usually, an atom, whatever its flavour, will have the same number of electrons as it has protons in its core. If an electron goes missing, or the atom collects an extra one, the positive and negative charges no longer balance each other out and the atom becomes what chemists call an ‘ion’ – a charged atom or molecule. Ions are important because their charges help stick together all sorts of substances, such as the sodium chloride of table salt and the calcium carbonate of limescale.

Besides making up kitchen cupboard ingredients, atoms form everything that crawls or breathes or puts down roots, building
stunningly complex molecules like DNA and the proteins that form our muscles, bones and hair. They do this by bonding (see page 20) with other atoms. What’s interesting about all life on Earth, however, is that despite its tremendous diversity, without exception it contains one particular flavour of atom: carbon.

From bacteria clinging to life around smoking hot vents in the deepest, darkest parts of the ocean to birds soaring in the sky above, there is not a living thing on the planet that doesn’t share that element, carbon, in common. But because we have not yet discovered life elsewhere, we can’t tell whether it was random chance that life evolved this way, or whether life could thrive using other types of atoms. Science-fiction fans will be well acquainted with alternative biologies – silicon-based beings have appeared in Star Trek and Star Wars as alien life forms.

Splitting the atom

J.J. Thomson’s early ‘plum pudding’ model of the atom viewed it as a doughy, positively charged ‘pudding’ with negatively charged ‘plums’ (electrons) distributed evenly throughout. That model has changed: we now know that protons and other subatomic particles called neutrons form the tiny, dense centre of the atom, and the electrons a cloud around them. We also know that protons and neutrons contain even smaller particles called quarks. Chemists don’t generally dwell on these smaller particles – they are the concern of physicists, who smash up atoms in particle accelerators to find them.

But it is important to remember that science’s model of the atom, and of how matter fits together in our Universe, is still evolving. The discovery of the Higgs boson in 2012, for example, confirmed the existence of a particle that physicists had already included in their model and used to make predictions about other particles. However, there’s still work to do to understand how the Higgs changes our view of the Universe.

Progress in the area of nanotechnology (see page 180), which promises everything from more efficient solar panels to drugs that seek and destroy cancer cells, has brought the world of the atom into sharper focus. The tools of nanotechnology operate at a scale of one billionth of a metre – still bigger than an atom, but at this scale it is possible to think about manipulating atoms and molecules individually. In 2013, IBM researchers made the world’s smallest stop-motion animation, featuring a boy playing with a ball. Both the boy and the ball were made from copper atoms, all visible individually in the movie. Finally, science is starting to work at a scale that matches the chemist’s view of our world.

The beauty of a living thing is not the atoms that go into it, but the way those atoms are put together.

Carl Sagan

The condensed idea

Building blocks

Chemists go to great lengths to discover new elements, the most basic chemical substances. The Periodic Table gives us a way to order their discoveries, but it’s not just a catalogue. Patterns in the table provide clues about the nature of each element and how they might behave when they encounter other elements.

The 17th-century alchemist Hennig Brand was a gold digger. After getting married, he left his job as an army officer and used his wife’s money to fund a search for the Philosopher’s Stone – a mystical substance or mineral that alchemists had been seeking for centuries. Legend had it that the stone could ‘transmute’ common metals like iron and lead into gold. When his first wife died, Brand found another and continued his search in much the same fashion. Apparently, it had occurred to him that the Philosopher’s Stone could be synthesized from bodily fluids, so Brand duly acquired no less than 1500 gallons of human urine from which to extract it. Finally, in 1669, he made an astounding discovery, but it wasn’t the stone. Through his experiments, which involved boiling and separating the urine, Brand had unwittingly become the first person to discover an element using chemical means.

Brand had produced a compound containing phosphorus, which he referred to as ‘cold fire’ because it glowed in the dark. But it took until the 1770s for phosphorus to be recognized as a new element. By this time, elements were being discovered left, right and centre, with chemists isolating oxygen, nitrogen, chlorine and manganese all within the space of one decade. In 1869, two centuries after Brand’s discovery, the Russian chemist Dmitri Mendeleev devised the Periodic Table and phosphorus assumed its rightful place therein, between silicon and sulfur.

For much of history, ‘the elements’ were considered to be fire, air, water and earth. A mysterious fifth, aether, was added to account for the stars, since they could not, as the philosopher Aristotle argued, be made from any of the earthly elements. The word ‘element’ comes from a Latin word (elementum) meaning ‘first principle’ or ‘most basic form’ – not a bad description, but it does leave us wondering about the difference between elements and atoms.

Decoding the Periodic Table

In the Periodic Table (see pages 204-5) elements are represented by letters. Some are obvious abbreviations, such as Si for silicon, while others, such as W for tungsten, seem to make no sense – these are often references to archaic names. The number above the letter is the mass number – the number of nucleons (protons and neutrons) in the nucleus of an element. The subscripted number is its number of protons (atomic number).

The difference is simple. Elements are substances, in any quantity; atoms are fundamental units. A solid lump of Brand’s phosphorus – incidentally, a toxic chemical and a component of nerve gas – is a collection of atoms of one particular element. Curiously though, not every lump of phosphorus will look the same, because its atoms can be arranged in different ways, changing the internal structure but also the outward appearance. Depending on how the atoms are arranged in phosphorus, it can look white, black, red or violet. These different varieties also behave differently, for instance, melting at wildly different temperatures. White phosphorus, for example, melts under the Sun on a very hot day, while black phosphorus would need to be heated in a roaring furnace at over 600 °C. Yet both are made from the very same atoms containing 15 protons and 15 electrons.

To the untrained eye, the Periodic Table (see pages 204–5) has the appearance of a slightly unorthodox game of Tetris, where – depending on the version you look at – some of the blocks have not quite dropped to the bottom. It looks like it needs a good tidy-up. Actually, it’s a well-ordered mess and any chemist will quickly be able to find what he or she is looking for among the apparent disarray. This is because Mendeleev’s cunning design contains hidden patterns that link together elements according to their atomic structures and chemical behaviours.

Along the table’s rows, from left to right, the elements are arranged in order of atomic number – the number of protons that each element has in its core. But the genius of Mendeleev’s invention was discerning when the properties of the elements began to repeat and a row should be turned. It is from the columns, therefore, that some of the more subtle insights are gleaned. Take the column on the far right, which runs from helium to radon. These are the noble gases, all colourless gases under normal conditions and all particularly lazy when it comes to being involved in any kind of chemical reaction. Neon, for instance, is so unreactive that it cannot be convinced to enter into a compound with any other element. The reason for this is related to its electrons. Within any atom, the electrons are arranged in concentric layers, or shells, which can only be occupied by a certain number of electrons. Once a shell is full, further electrons must start to fill a newer, outer layer. Since the number of electrons in any given element increases with increasing atomic number, each element has a different electron configuration. The key feature of the noble gases is that all of their outermost shells are full. This full structure is very stable, meaning the electrons are difficult to prod into action.

The world of chemical reactions is like a stage ... The actors on it are the elements.

Clemens Alexander Winkler, discoverer of the element germanium

We can recognize many other patterns in the Periodic Table. It takes more effort (energy) to prise an electron away from an atom of each element as you move from left to right, towards the noble gases, and from bottom to top. The middle of the table is occupied mostly by metals, which become more metallic the closer you edge to the far left corner. Chemists use their understanding of these patterns to predict how elements will behave in reactions.

One of the few things that chemistry shares in common with boxing is that both have their superheavyweights. While the flyweights float at the top of the Periodic Table – with atoms of hydrogen and helium carrying just three protons between them – those on the bottom rows have sunk by virtue of their heavy atomic loads. The table has grown over many years to incorporate new discoveries and heavier elements. But at number 92, the radioactive element uranium is the last element found in nature. Although the natural decay of uranium yields plutonium, the quantities are vanishingly small. Plutonium was discovered in a nuclear reactor and other superheavyweights are made by smashing together atoms in particle accelerators. The hunt isn’t over yet but it’s certainly become a lot more complicated than boiling up bodily fluids.

The hunt for the heaviest superheavy

No one likes a cheat, but you’ll find them in every profession and science is no exception. In 1999, scientists at the Lawrence Berkeley Laboratory in California had published a scientific paper celebrating their discovery of superheavy elements 116 (livermorium) and 118 (oganesson). But something wasn’t making sense. Having read the paper, other scientists had tried to repeat the experiments, but no matter what they did they couldn’t seem to conjure a single atom of 116. It turned out one of the ‘discoverers’ had fabricated the data, leaving a US government agency to make an embarrassing climbdown from statements about the world-class science it was funding. The paper was pulled and the plaudits for discovering livermorium went to a Russian group a year later. The scientist who faked the original data was fired. Such is the prestige associated with discovering a new element these days that scientists are willing to stake their entire careers on it.

The condensed idea

The simplest substances

Isotopes aren’t just deadly substances used to make bombs and poison people. The concept of an isotope is one that encompasses many chemical elements that have a slightly altered quota of subatomic particles. Isotopes are present in the air we breathe and the water we drink. You can even use them (perfectly safely) to make ice sink.

Ice floats. Except when it doesn’t. Just as all atoms of a single element are the same, except when they are different. If we take the simplest element, hydrogen, we can agree that all atoms of this element have one proton and one electron. You couldn’t call a hydrogen atom a hydrogen atom unless it had only one proton in its nucleus. But what if the single proton was joined by a neutron? Would it still be hydrogen?

Neutrons were the missing piece of the puzzle that eluded chemists and physicists until the 1930s (see The missing neutrons, opposite). These neutral particles make no difference at all to the overall balance of charge in an atom, but radically alter its mass. The difference between one and two neutrons in the core of a hydrogen atom is enough to make ice sink.

Packing an extra neutron into a hydrogen atom makes a big difference – for these flyweight atoms, it’s double the quota of nucleons. The resulting ‘heavy hydrogen’ is called deuterium (D or 2H) and, just as regular hydrogen atoms do, deuterium atoms hook up with oxygen to make water. Of course, they don’t make regular water (H2O), but water with extra neutrons in it: ‘heavy water’ (D2O), or to give it its proper name, deuterium oxide. Take heavy water – easily purchased online – and freeze it in an ice-cube tray. Plop a cube into a glass of ordinary water and, bingo, it sinks! For comparison, you can add an ordinary ice cube and marvel at the difference that one subatomic particle per atom makes.

In nature, about one in every 6,400 hydrogen atoms have an extra neutron. There is, though, a third type – or isotope – of hydrogen, and this one is much rarer and rather less safe to handle at home. Tritium is an isotope of hydrogen in which each atom contains one proton and two neutrons. Tritium is unstable, however, and like other radioactive elements it undergoes radioactive decay. It is used in the mechanism that triggers hydrogen bombs.

Often the word ‘isotope’ is preceded by the word ‘radioactive’, so there might be a tendency to assume that all isotopes are radioactive. They are not. As we have just seen, it is perfectly possible to have an isotope of hydrogen that is non-radioactive – in other words, a stable isotope. Likewise, there are stable isotopes of carbon, oxygen and other elements in nature.

Unstable, radioactive isotopes decay, meaning that their atoms disintegrate, shedding matter from their core in the form of protons, neutrons and electrons (see Types of radiation, on page 14). The result is that their atomic number changes and they can become different elements altogether. This would have seemed like magic to 16th- and 17th-century alchemists who were obsessed with finding ways of changing one element into another (the other, ideally, being gold).

The missing neutrons

The discovery of neutrons by physicist James Chadwick – who went on to work on the atomic bomb – solved a niggling problem with the weights of the elements. For years, it had. . .