A Little History Of Science: Tabling the Elements
Every time we mix ingredients to bake something, we are using chemical reactions. The fizzing as we descale our kettles is chemistry at work for us. The plastic water bottles we carry, the coloured clothes we wear, are possible because of chemical knowledge gained over hundreds of years.
Chemistry became modern in the nineteenth century. Let’s recap a little. At the beginning of the century, chemists embraced Dalton’s original idea of the atom, as you read in Chapter 21. Then they made great strides in creating a special language that they would all understand, whatever country they came from. They had the system of symbols for elements, such as H2 for two atoms of hydrogen. Everyone agreed that the atom was the smallest unit of matter. They used the word element for a substance made of only one kind of atom (carbon, for example). A compound was two or more elements bonded together chemically. You could break down compounds into elements (ammonium could be broken down into nitrogen and hydrogen), but once you got to the individual elements, you couldn’t break things down any further.
Although atoms were clearly not the hard tiny balls that Dalton had suggested, it was extremely difficult to say exactly what they were. Instead, chemists began to discover a lot about how they acted when placed with other atoms or compounds. Some elements were simply not interested in reacting with others, no matter what you did. Some would react so violently together that you had to guard against an explosion. Sometimes, however, you could get a reaction if you helped to get it started. Oxygen and hydrogen could be placed together in a flask and nothing happened. If you put a spark to it, you had to watch out! Despite the dramatic explosion, the reaction produced nothing more unusual than water. At the other extreme, if magnesium and carbon were put together in a flask with no air, you could heat them forever and nothing would happen. Let in a bit of air, and you would be greeted with bright light and an awful lot of heat.
Chemists were aware of these various chemical reactions. They also became curious about what caused them and the patterns revealed in the laboratory. They set about their experiments in two main ways: synthesis and analysis. Synthesis is putting elements together: you start with single elements or simple compounds, and when these react with each other, you look at the results – at what has been made. Analysis is the reverse: you start with the more complex compound, do something to break it down, and, by looking at the end products, try to understand the compound that you started with. These methods began to give chemists a good idea of what many fairly simple compounds consisted of. It was then easier to create more complicated compounds, by adding new bits to substances they had a fair idea about.
From all these experiments, two things became particularly clear. First, as we have seen, the elements themselves each seemed to have either positive or negative tendencies. As the old saying goes, opposites attract. For example, sodium, a naturally positive element, combined easily with chlorine, a negative one, to make sodium chloride (which is just the table salt we sprinkle over food). The positive and negative cancel each other out, so salt is neutral.
All stable compounds (those that won’t change unless something is done to them) are neutral even though they are made up of elements that were not necessarily so. Combining sodium and chlorine is an example of synthesis. You can do chemical analysis of the salt you’ve made. Dissolve the salt in water, put the solution in an electric field with its positive and negative poles, and it will split up. Sodium migrates to the negative pole, chlorine dances to the positive one. Hundreds of similar experiments convinced the chemists that the atoms of such elements do indeed have these positive and negative characteristics. And these characteristics play an essential role in determining what happens when elements react with each other.
Second, some groups of atoms may stick together during experiments, and these collective atoms can act like a single unit. These units were called ‘radicals’ and they too are positive or negative.
They were especially important in ‘organic’ chemistry, where chemists were coming to understand a whole series of related compounds (all of which contained carbon), such as ethers, alcohols or benzenes. Benzenes were a fascinating group, each with a ring-like structure. Many chemists were eager to try to classify these organic groups, to understand what they were made of and how they reacted – not least because a lot of the substances were becoming valuable to industry. Increasingly, such industrial chemicals were made not in small laboratories, but in factories. The demand was growing for fertilisers, paints, medicines, dyes and, especially from the 1850s, oil products. The modern chemical industry had begun, and chemistry became a career, not just an indulgence for the curious or the rich.
The elements, too, have their own unique chemical and physical properties. As more and more were discovered, chemists found certain patterns. It appeared that individual atoms of some elements, such as hydrogen, sodium or chlorine, only wanted to combine with other atoms singly. For example, a single atom of hydrogen and one of chlorine combined to make a powerful acid, hydrochloric acid (HCl). A single atom of others, such as oxygen, barium and magnesium, seemed to have a double capacity to combine with other atoms or radicals, and so it takes two atoms of hydrogen to combine with oxygen to make water. Some elements were even more flexible, and there were always exceptions that made any hard and fast rules difficult to set down. Elements (and radicals) also differed in their eagerness to enter into chemical reactions. Phosphorous was so active that it had to be treated carefully; silicon was generally sluggish and much less dangerous.
The elements differed dramatically in their physical properties, too. At normal temperatures, hydrogen, oxygen, nitrogen and chlorine were gases; mercury and sodium were liquids. Most were naturally solids: metals like lead, copper, nickel and gold. Many other elements, above all carbon and sulphur, both intensively studied, were normally in a solid state. Put most solids in an ordinary furnace and they can easily be melted, and sometimes even vaporised (turned into a gas). Liquid mercury and sodium were also easy (if dangerous) to vaporise. Nineteenth-century chemists were not able to get low enough temperatures to turn gases like oxygen and nitrogen into liquids, much less solids. But they recognised their problem was a merely technical one. In principle, each element could exist in each of the three states of matter: solid, liquid and gas.
By the 1850s, chemistry was coming of age, and in this exciting period there was much to debate, about the relative weights of atoms, how molecules (groups of atoms) were bound together, the differences between ‘organic’ and ‘inorganic’ compounds, and much else. In 1860, something happened that helped create modern chemistry. It was something that today seems quite ordinary, but was unusual then: an international meeting. In the days before telephone, emails and easy travel, scientists rarely met and they communicated mostly by letters. Hearing another scientist from abroad talk about their work, with an open discussion afterwards, was a rare event. International meetings began in the 1850s, helped by more available travel by train and steamship, and they allowed people to meet and talk with their colleagues from other countries.
They also announced to the world a belief widely shared by the scientific community: that science itself was objective and international, and above religion and politics, which often divided people and set whole nations at war with one another.
The 1860 chemistry gathering met for three days in Karlsruhe, Germany. Many of the leading young chemists from all over Europe came there, including three who would direct chemistry for the rest of the century. The meeting’s aims were set by the German August Kekulé (1829–96). He wanted chemists from different countries to agree on the words they should use to define the substances they worked with, and the nature of atoms and molecules. A fiery Sicilian chemist, Stanislao Cannizzaro (1826– 1910), had already been arguing for this, and he gladly participated. So did a passionate Russian chemist, from Siberia, Dmitry Ivanovich Mendeleev (1834–1907). The delegates discussed Kekulé’s suggestions for three days, and while no complete agreement was reached, the seeds had been sown.
At the meeting, copies of an article published by Cannizzaro in 1858 were given to many of the delegates. Here he reviewed the history of chemistry during the earlier part of the century. He called for chemists to take seriously the work of his fellow countryman Avogadro, who had clearly distinguished between an atom and a molecule. Cannizzaro also argued that it was vital to deter- mine the relative atomic weights of the elements, and he showed how this could be done.
Mendeleev got the message. He owed much to his formidable mother, who had taken this last of her fourteen children from Siberia to St Petersburg, so that Mendeleev could learn about chemistry properly. Like many outstanding chemists of the time, Mendeleev wrote a textbook, based on his own experiments and what he taught his students. Like Cannizzaro, he wanted to bring order into the many elements that had been identified. Patterns had already been revealed: what were called the ‘halogen’ family – chlorine, bromine and iodine, for instance – reacted in similar ways. They could also be swapped for each other in chemical reac- tions. Some metals, such as copper and silver, also shared similarities in their reactions. Mendeleev began to list the elements in the order of their relative atomic weight (still using hydrogen as ‘1’).
He presented his ideas in 1869.
Mendeleev did more than simply compile a list of the elements by atomic weight. He created a table, with rows and columns. You could read it across as well as up and down, and could see the relationship between elements with similar chemical properties. At first, his periodic table, as he came to call it, was very rough, and few chemists paid much attention to it. As he began to fill in the details, something interesting happened: there seemed to be occasional missing elements here and there, substances that his table implied should be there, but which had not been discovered. In fact, there was a whole missing column in his table, predicted by the relative atomic weights. Years later, this column turned out to be filled by non-reactive gases – called the ‘noble’ gases. Like aristocratic noblemen who don’t mix socially with people below them, these gases were aloof from chemical reactions. The main ones were discovered only in the 1890s, and Mendeleev did not accept the findings at first. He soon realised that helium, neon and argon, with the atomic weights that they were shown to possess, had been predicted by his periodic table.
In the 1870s and 1880s, chemists discovered several more of the elements that Mendeleev had predicted on the basis of his table. Many chemists had dismissed as crazy speculation his predictions that the elements eventually called beryllium and gallium must exist. As the gaps he had identified gradually began to be filled, chemists appreciated the power of Mendeleev’s table. It was guiding them to discover new elements in nature. It was also explaining what each element is like and how it reacts with other chemicals.
What began as Mendeleev’s attempt simply to understand the elements produced an amazing key to how nature behaves. The periodic table now hangs in classrooms and chemical laboratories all over the world. For much of the nineteenth century, chemists had been concerned with chemical composition: which atoms and radicals made up specific compounds. The brains behind that first international chemical congress, August Kekulé, began to go further. He encouraged scientists to aim to understand chemical structure.
Today’s chemistry and molecular biology depend on scientists knowing how atoms and molecules are arranged in a substance: where they all sit, and the shapes they form. It would be impossible to search for new drugs without this understanding, and Kekulé pioneered it. He told of a dream in which he saw a chain of carbon atoms curled around itself, like a snake biting its own tail. This inspired one of his greatest insights, into benzene, the compound of hydrogen and carbon, which has a closed ring structure. Radicals or elements can be added at various points around the ring, and this was an important advance for organic chemistry.
Dreams are one thing. Doing the hard slog is another. Kekulé spent many hours in his laboratory, experimenting. He made sense of organic chemistry – the chemistry of carbon compounds – and taught the whole chemical world how to classify them in their natural families. He was fascinated by carbon’s flexibility in joining with other chemicals. Methane gas, then widely used for heat and light, was CH4 – one carbon atom joined to four hydrogen atoms.
Two oxygen atoms could combine with a carbon atom, giving CO2, carbon dioxide. That these atomic preferences were not set in stone was shown by the fact that carbon and oxygen could combine as single atoms to create CO, the deadly gas carbon monoxide.
Chemists came up with a word for these joining patterns: valence. And it could be deduced from the position of each element in Mendeleev’s periodic table. They speculated about why this was so. Real understanding came only with the discovery, by the physicists, of the inner structure of the atom, and of the electron. The electron linked the chemist’s atom with the atom that the physicists were studying, and the next chapter will tell that story.