It all started with alchemy. Of course, the ancient world had had some practical knowledge of chemistry and had seen some advances in the applied chemistry long before the emergence of alchemy. The distant roots of chemistry are probably to be found in Bronze Age (roughly 3 300 BCE to 1 200 BCE) Mesopotamian and Egyptian medicine.
However, it is not until the emergence of the Classical Greek civilisation (during the late sixth century BCE) that a modern conception of science first emerged – that is, a conception that natural, and not divine, causes could be found for natural phenomena.
“The Greek achievement in chemistry presents us with an odd contrast,” observes JG Landels. “On the one hand, Greek doctors knew and used a very wide range of chemi- cals, both mineral substances and others extracted from animals and plants by various processes. But, on the other, there was virtually no general theory about chemicals, or the laws which govern their combination or separation. However, it is as well to recall that the Greeks and Romans managed their practical chemistry with fair success.” (Engineering in the Ancient World, revised edition, Constable, London, 2000, pp 197 to 198.)
Alchemy to Chemistry
Make no mistake, alchemy was a pseudoscience. It certainly represented a regression on the early scientific conceptions of ancient Greece and Rome, because alchemy contained a large dose of supernatural belief.
“Beginning about the year 100 CE and reaching its flower in medieval times, alchemy was an art based partly upon experimentation and partly upon magic,” points out James Randi. “Early investigators of natural processes centred their research on a mythical substance they knew as the philosopher’s stone . . . which was supposed to possess many valuable attributes, such as the power to heal, to prolong life, and to change base metals into precious metal – such as gold . . . but along the way, alchemists made many genuinely valuable contributions to knowledge, though such fundamental discoveries as the chemical elements and the manner in which they form compound substances escaped them . . . eventually, when the nonsense and misinformation were boiled out of alchemy, it became chemistry.” (https://web.randi.org/a---encyclopedia-of-claims.html, accessed February 15, 2019.)
Not until the seventeenth century did modern chemistry begin to emerge. A key step in the extraction of the science from the previous pseudoscience of alchemy was the publication, in 1661, of the book The Sceptical Chymist, written by Robert Boyle (1627–1691), although Boyle did not represent a clear break with alchemy. While he stripped the super- natural away from alchemy, he still believed that the fundamental alchemical goal of transmutation (transforming base metals into precious ones) was indeed possible, but using purely scientific means.
Chemistry was fully established as a proper science, completely divorced from alchemy, in the eighteenth century. The central figure in this achievement was Frenchman Antoine-Laurent Lavoisier (1743–1794). He firmly established the quantitative scientific approach to chemistry, and systematically determined the weights of reagents and products, including gaseous ones, involved in chemical reactions. He defined the ‘element’ in chemistry as a substance that chemical analy- ses could not break down into simpler components. He invented the modern system of chemical nomenclature – he was responsible for naming carbon, hydrogen and oxygen. He established oxygen’s role in both respiration and combustion. He showed that water was made up of hydrogen and oxygen, and that diamond was a form of carbon. His book, Elements of Chemistry, is regarded as the first modern chemistry textbook.
As a result of all his work, Lavoisier produced a list of 33 elements. These included hydrogen, mercury, nitrogen, oxygen, phosphorous, sulphur and zinc. However, the list also included the nonexistent ‘caloric’ and light, a reminder that his work marked the start of modern chemistry, not its perfection.
From List to Table
After Lavoisier, steady progress was made. The first laws of chemistry were discovered. The role of atoms in the formation of chemical compounds was realised. Research into establishing the weight of atoms (the modern term is relative atomic mass) became a subject of great importance. The list of known elements grew longer, as researchers discovered more of them. They began to become aware of the existence of patterns in the properties of various elements. This led to some chemists dividing the overall list of elements into sub-lists of elements with similar characteristics. Key figures in this process included Johann Döbereiner (who published his work in 1829), Alexandre-Emile Béguyer de Chancoutois (research published in 1862) and John Newlands (who, in 1863, divided the then known 56 elements into eleven groups, based on their atomic structure).
The scene was finally set. The breakthrough was made by Russian chemist Dmitri Mendeleev (1834–1907). The foundation of his achievement was the advances in research into atomic weights (relative atomic masses), in which field the names of the Italians Amadeo Avogadro and Stanislau Cannizzaro are particularly prominent.
Mendeleev’s stimulus was that, having been appointed professor of general chemistry by the University of St Petersburg in 1867, he discovered that there was no satisfactory Russian textbook on inorganic chemistry, so he had to write one. Having written the first volume, he had difficulty establishing a framework for the second volume. He later explained, in his 1905 work, Principles of Chemistry, what he did. “So I began to look about and write down the elements with their atomic weights and typical properties, analogous elements, and like atomic weights on separate cards, and this soon convinced me that the properties of the elements are in periodic dependence upon their atomic weights.”
A man who whiled away the time on long train journeys across Russia by playing solitaire, or patience, Mendeleev arranged his elements cards in vertical columns and horizontal rows, like playing cards in these games. The then known elements were arranged in the columns in the order of increasing atomic weight (relative atomic mass). New columns were created when this allowed him to place elements with similar characteristics in the same row. This prototype periodic table was published in Russia in 1869. He subsequently revised and refined it.
Now, it only became possible to measure atomic mass after invention of the mass spectrometer in 1912 (by physicist JJ Thomson). However, in 1811, Avogadro realised that the volume of a gas, at a given temperature and pressure, was proportional to the number of atoms or molecules of which it was composed; this applied to all gases. The Carbon-12 atom became the reference unit, and the atomic weights of other atoms were estimated relative to Carbon-12. This involved varying degrees of educated guesswork on the part of chemists.
As a result of this situation, when Mendeleev found an element whose atomic weight indicated that it should be in one place on his table, but all its other characteristics indicated it should be in another place, he did not hesitate to unilaterally change that element’s atomic weight to bring it into the place that its other characteristics pointed to. Most of his changes were subsequently vindicated.
Perhaps most importantly, he left spaces in his table for elements that had not yet been discovered. This had never been done before. Not only that but, using the patterns of characteristics revealed by his table, Mendeleev actually predicted the properties of five then undiscovered elements and their compounds. Three of these were discovered by 1886 (namely gallium in 1875, scandium in 1879 and germanium in 1886), and his predictions were proven to have been extremely accurate. The reputation of the Periodic Table was firmly established.
There was one last glitch that had to be overcome. In the 1890s, British chemist William Ramsay and his coworkers discovered the noble, or inert, gases, starting with argon in 1894, followed by neon, krypton and xenon, all in 1898. Further, in 1895, he established that helium (discovered in the sun in 1868 using the then new technique of spectroscopy) existed on earth. At first, these did not seem to fit in with Mendeleev’s Periodic Table, which had placed the elements in seven groups. Then Ramsay realised that the noble gases together represented an eighth group.
“Others before [Mendeleev] had suggested that the list of known elements might be arranged in a meaningful pattern,” pointed out British science historian Mike Sutton in the Royal Society of Chemistry journal Chemistry World. “They noted significant correspondences, but found no definitive picture. Mendeleev, however, was convinced that the chemical elements had to be viewed as a collective entity. Armed with this conviction, he gave his table coherence by boldly revising the positions of some known elements, and by leaving gaps for others yet undiscovered. Although some of his predictions were incorrect, he scored enough hits to establish his table as the basis for our understanding of the elements, and to confirm his status as one of the founders of modern chemistry.”
The Periodic Table Today
Following the discovery of radioactivity in 1896, it came to be realised that the characteristics of an element were determined not by its atomic weight but by its atomic number, which is the number of protons in the nucleus of an atom. The Periodic Table thereafter became based on atomic number instead of atomic weight, a development led by British chemist Henry Moseley. Amazingly, this change did not alter the table’s arrangement.
As new elements were discovered, they were added to the Periodic Table. For reasons of publishing convenience, two series of elements, the lanthanides and the actinides, are presented as two separate bands, located below the main body of the Periodic Table. This is simply because incorporating them into the main body – which would be scientifi- cally appropriate – would make the Periodic Table too long to conveniently publish on a page of a book.
The lanthanides number 15 elements, so called because they are chemically similar to lanthanum. Although cerium was discovered in 1805, lanthanum itself in 1839 and terbium in 1843, most of the lanthanides were discovered and named in the period 1879 to 1907.
The actinides also number 15 elements, again grouped together because they have similar chemical properties, and all of them radioactive. They are named after the first element in the series, actinium (discovered in 1899), but the most famous and most common actinide element is uranium (discovered in 1789), while the second-most famous is plutonium (discovered at the end of 1940). Only five of the actinides occur in nature; the rest are human made in laboratories, following the invention of the cyclotron by American physicist Ernest Lawrence. Many of the synthetic actinides were discovered at the University of California, in Berkeley, in the 1940s and 1950s, under the leadership of two US scientists: chemist Glenn Seaborg and physicist Albert Ghiorso. One of these human-made elements, first synthesised in 1955, which has an atomic number of 101, was named mendelevium in honour of Mendeleev.
Since then, more and more elements have been synthesised in laboratories around the world and added to the Periodic Table. Most of them have incredibly brief half-lives, before they decay into other elements. There are now 118 known, confirmed elements, of which the four most recently discovered (all during this century) are moscovium (atomic number 115), livermorium (116), tennessine (117) and oganesson (118).
But what is the information contained in the Periodic Table? Firstly, it is arranged from left to right, and top to bottom, in terms of increasing atomic number. The horizontal bands are called periods; the vertical columns are called groups. The elements in each group display strong similarities in their chemical properties, while a gradation in these properties may be observed along each of the periods.
The elements in the Periodic Table are divided into metals, nonmetals, and metalloids (which have some of the characteristics of metals and some of the characteristics of nonmetals – silicon is an example). Some of the metals are further categorised as transition metals or transitional elements. This term refers to their electronic structure – that is, which of their electron ‘orbitals’ are filled.
The groups containing the nontransitional elements are numbered I to VII, and 0. These numbers refer to the number of electrons in the outermost, unfilled shells of these elements. (The atomic nucleus is surrounded by ‘shells’ containing electrons; these can be subdivided into subshells, each of which encompasses one or more orbitals, each of which, to simplify, can be thought of as containing an electron.) This number determines the maximum valency of the element – that is, its ability to combine with other elements. The group numbered 0 is composed of the noble elements, which very rarely combine with other elements.
Each square making up the Periodic Table contains key information on one of the elements. There is the element’s name – for example, tantalum – its symbol (in this case, Ta), as well as, at the top, its atomic number (for tantalum, 73) and, at the bottom, its relative atomic mass (in the case of tantalum, 180.9479).
Once it is understood, the Periodic Table serves a kind of information-filled one-stop, easily accessible visual filing cabinet. Little wonder chemists value it so much. And little wonder that it has also become a symbol – dare one say icon – of the science of chemistry itself.