[This series is a companion to our International Year of the Periodic Table project — Timeline of Elements and Mendeleev Mosaic. Time periods from our timeline project will be featured highlighting the historical and scientific advancements resulting in the discoveries of the elements of that period. Jim Marshall researches the history of the discovery of the elements. His studies have carried him to over 20 countries where he has photographed and documented specific sites which he describes in his publications on chemical history. Explore the artwork for the Timeline of Elements for Davy's Elements with individual discovery stories of each element with acknowledgement to artist(s), teacher, school and countries on our Timeline of Elements website.]
Historically, the distinction between sodium and potassium was not made easily. Only gradually was it recognized that there was a difference between “plant alkali” (potash, potassium carbonate), obtained from the ashes of plants, and “mineral alkali” (soda, sodium carbonate), found in salt flats and evaporated seawater. In 1683 Johnan Bohn prepared “cubic saltpetre” (sodium nitrate) from a mixture of salt and nitric acid, which he recognized as different from the ordinary saltpetre (potassium nitrate) by the shape of its crystals and its taste. Stahl (the promoter of phlogiston; see previous chapter on the “True Elements”) formally distinguished between “natural alkali” (soda) and “artificial alkali” (potash) in 1702. Several other investigators continued to characterize the two alkalis, with the most definitive account by Marggraf in 1759 who prepared very pure samples of “cubic saltpetre” and “prismatic saltpetre” from which he made gunpowder (from sulfur, charcoal and saltpetre). The “cubic saltpetre gunpowder” flashed yellow and the “prismatic saltpetre gunpowder” flashed blue. But the true nature of the two alkalis remained unknown.
With Lavoisier’s amazing insight (see the “True Elements”), it may seem strange that he did not include sodium and potassium in his list of the elements. He explained that he did not include the “fixed alkalis” in his
list of thirty-three substances because he believed they were compounds, although the “nature of the principles which enter into their composition is still unknown.” He suspected that they might contain nitrogen, since it had been shown that ammonium contained that element. It remained for Sir Humphry Davy at the beginning of the nineteenth century, to decompose soda and potash to the elemental substances.
Davy was well educated and became an assistant lecturer and director of the laboratory at the Royal Institution at London. He became interested in electrochemistry and tried to decompose the caustic alkalis with a voltaic pile (the “caustic alkalis” are sodium hydroxide and potassium hydroxide, in contrast to the “mild alkalis” which are the respective carbonates). First attempts accomplished nothing but decomposing the water of the aqueous solutions. Recognizing his mistake, he fused anhydrous potash itself, and to his great delight observed flames at the cathode — he is reported to have danced about the laboratory with joy when he observed globules, appearing like quicksilver, being generated at this negative pole. He was
able to collect a few of these globules, which he tossed into water; they raced about swiftly and burst into flame with a beautiful lavender light. Davy continued his investigation and found the flames were caused by the ignition of hydrogen which was being generated upon contact with moisture. He named this element potassium after potash. Davy immediately turned to soda, but found it could be generated only with “greater intensity of action in the batteries,” (i.e., greater voltage). He named this new element sodium after soda, only a few days after the discovery of potassium. When tossed into water, it burst into yellow flame.
Many chemists believed that these “elemental metals” sodium and potassium were in fact compounds made up of the alkalis + hydrogen. This hypothesis appeared to be reasonable, since ammonium (“volatile alkali,” i.e., ammonium hydroxide) was ammonia + hydrogen. However, it was promptly shown that no hydrogen could be evolved from the metal itself. It was realized the hydrogen must be generated from the water: K (or Na) + H2O → KOH (or NaOH) + H2. The elemental nature of metallic sodium and potassium was thus established.
Flushed with success, Davy turned to the alkaline earths. Compounds of calcium had figured prominently in the technology of the ancients. The ancient Romans understood that limestone (calcium carbonate) when heated in a kiln produced slaked lime (calcium oxide), which served as excellent mortar. The explanation was that “water had boiled out of the limestone,” not correct but consistent with the observation that weight was lost in the process. Another compound of calcium, gypsum (calcium sulfate) when heated produced another form of mortar used in Egypt. This mortar, which had lost water, was later called Plaster of Paris. Later explanations of this slaking of lime (barium hydroxide) included the interpretation that slaked lime and baryta were elements, although Lavoisier correctly believed they were oxides of elements not yet prepared. Methods then available — principally heating with charcoal — were employed to provide the metallic elements by reduction of the oxides. These methods, however, were not powerful enough, and the preparation of the metallic elements had to await the electrolysis technique of Davy.
In the 1750s, the work of Joseph Black, a professor of chemistry at Edinburgh, Scotland, was pivotal in elucidating the true nature of the slaking (calcining) process. Black investigated both lime and magnesia alba, a medicine sold in Rome. He determined that the loss of weight during the calcining process was due to a compound he called “fixed air” (carbon dioxide). This air could be “fixed” (made solid) by reacting again with the metal oxide. Working principally with calcium and magnesium compounds, Black showed that:
limestone → quicklime + fixed air,
i.e., CaCO3 → CaO + CO2
magnesia alba → calcined magnesia + fixed air,
i.e., MgCO3 → MgO + CO2
limestone + acid → calcium salt + water + fixed air,
i.e., CaCO3 + 2H+ → Ca2+ + H2O + CO2
magnesia alba + acid → magnesia salt + water + fixed air,
i.e., MgCO3 + 2H+ → Mg2+ + H2O + CO2
Black came close to suggesting the atomic theory of gases by suggesting that the fixed air was dispersed through the atmosphere in the form of an “extremely subtile powder.” Black continued his studies to show that fermentation and combustion evolve fixed air by mixing this air with lime water, producing the milky product of precipitated lime (calcium carbonate). This test of identifying carbon dioxide exists to this day in the laboratory: CaO + CO2 → CaCO3.
Yet another alkaline earth oxide — baryta — had been produced by Scheele in 1774. This earth, named because of its heaviness, had been first observed in “Bologna stone,” discovered in the 1600s by Vincenzo Casciarolo, a shoemaker and alchemist in Bologna. This stone phosphoresced when heated.
Meanwhile, a mineral from Strontian, Argyleshire, Scotland was analyzed by Crawford and Hope, both of whom concluded in the late 1700s that an earth was present that was intermediate between calcium and barium.
The stage was now set for Davy to demonstrate that his method for potassium and sodium could be used as a general procedure. One by one he decomposed salts of calcium, magnesium, strontium, and barium by a special recipe he developed (as suggested by Berzelius, who had already experimented with voltaic cells): he electrolyzed a mixture of moist alkaline earth oxide and mercury oxide to produce the amalgamated metal. The free metal was then prepared by volatilizing the mercury by heating. Davy produced each of these metals in 1808.
The sample of magnesium that Davy prepared was very small — it was not until 1831 that Bussy prepared magnesium in substantial amounts not by electrolyzing the salts but by reacting the salts with metallic potassium.
Another element that caused Davy difficulty was lithium. Lithium had been known only for a short period of time — the presence of this element was not suspected until the early 1800s when Arfwedson (in Berzelius’ laboratory) analyzed petalite (LiAlSi4O10) and found a difference between the observed and calculated total weight. Too small to be accounted for by sodium, the discrepancy was explained by the presence of a new, lighter alkali (in 1818). Several investigators, including Davy, promptly tried to decompose the alkali salt by electrolysis, but only minute and impure samples could be obtained. Only in 1855 could sufficient quantities of pure lithium be prepared, from pure lithium chloride by Bunsen and Matthiessen in Germany using a powerful battery bank. Today one can understand the difficulty of reducing lithium salts into lithium metal when it is realized that lithium is the most electropositive element in the electrochemical series (hence, its powerful nature in today’s batteries).
Finally, Davy was able to prepare boron in elemental state; previously, it was known only in the form of borate salts. The compound that Davy used in this preparation was boric acid (“sedative salt”), H3BO3. This element is technically not a metal, but a metalloid (properties between those of a metal and a nonmetal). Elemental boron was simultaneously being prepared by the French chemists Louis Jacques Thénard and Joseph Louis Gay-Lussac.
Davy's insight and major contribution to the elements was his proof that two distinctly different groups of chemically reactive substances — the alkali/alkaline metals and the halogens, were not compounds but instead true elements. For example, chlorine was generated by Scheele (see March 2019 issue) but it was not realized as an element until Davy's investigations. Similarly, the French chemists did not realize that iodine was an element until Davy showed that passing iodine over a red-hot filament did not alter it. Iodine was discovered by Bernard Courtois at a nitrary outside Paris while Davy was conducting his research in London (a “nitrary was a country farm supporting the French war effort by producing saltpeter (potassium nitrate) for gunpowder from rotting heaps of manure, garbage, and urine). Courtois observed violet fumes from varec (seaweed, the source of potassium) when reacted with sulfuric acid, which he subsequently collected as crystals of purple iodine.
Although primarily known for his work with the reactive metals, Davy also promoted the use of nitrous oxide N2O (laughing gas) as an anesthetic during simple surgical procedures, such as in dentistry. Davy also produced a “safety lamp” for the tin miners of Cornwall, which was his original home. A statue has been erected in his honor which stands in Penzance, his home town in western Cornwall.
Davy was a captivating lecturer and was responsible for the success of the famous lecture series at the Royal Institution in London. His first lecture was in 1801; his protégé Michael Faraday inaugurated the Christmas Lectures in 1825 which survive to this day and are a famous attraction to the international scientific community.
Photos taken from J.L. Marshall, Discovery of the Elements, Peterson Custom Publishing, 2002.