Atoms & Light

Part 6 - Periodic Tables

In the late 1700s, the French chemist Joseph Proust noted that oxygen always reacted with other elements in fixed proportions. One part oxygen always reacted with two parts of hydrogen to form water and one part oxygen always reacted with to one part mercury to form mercury calx. Some substances could combine to form new compounds while others could be separated to form simpler ones, and some (elements) could not be broken down any further. 

In 1803 Britain, John Dalton, observing the morning fog, realized that water could exist as a gas that mixed with air but ice could not mix with air. So why would liquid water sometimes behave as a solid and sometimes as a gas? Dalton hypothesized that all matter was composed of tiny particles. In the gas state, those particles floated freely around and could mix with other gases, as Bernoulli had proposed. But Dalton extended this idea to all matter – gases, liquids and solids. He refined these concepts in his 1808 paper, A New System of Chemical Philosophy. It was the first modern atomic theory.

Dalton proposed that all matter was composed of indivisible and indestructible atoms, tiny particles in various states of motion, tightly packed together in solids with no room to move, less tightly packed in liquids so that they could move relative to one another and even less densely packed in gases and free to move randomly in any direction. Solids had a definite shape and volume. Liquids had a definite volume but no definite shape. Gases had no definite volume and no definite shape.

All atoms of each element were identical; atoms of different elements having different properties. Every atom of an element was identical to every other atom of the same element; atoms of different elements, such as oxygen and mercury, were different from each other and could be characterized according to their atomic weight. As the absolute weights of atoms could not be determined at that time, early measurements were made by comparing weights of various atoms to hydrogen which was assigned the number one. Scientists noted that many elements were not exactly multiples of one, but it was not until the discovery of isotopes in the late 1800's that the discrepancies would be explained. Chemical reactions involved the combination of atoms, not the destruction of atoms. Atoms were unchangeable, so compounds, such as water and mercury calx, were formed when one atom chemically combined with one or more atoms of other elements.

Elements combined to form compounds in whole-number ratios according to precise formulas. Water, for example, was always made up of two parts hydrogen and one part oxygen. (A single "particle" of a compound made up of two or more atoms is called a molecule).

In 1829, Johann Wolfgang Döbereiner noted that, based on their chemical properties, many of the elements could be grouped together. Lithium, sodium and potassium, for example, were all soft, reactive metals. In 1857, German chemist August Kekulé noticed that carbon often had four other atoms bonded to it. Methane, for example, had one carbon atom and four hydrogen atoms. This concept eventually became known as valency.

In 1862, the French geologist Alexandre-Émile Béguyer de Chancourtois noticed the periodicity of the elements. When the elements were arranged in a helix on a cylinder by order of increasing atomic weight, he showed that elements with similar properties seemed to occur at regular intervals.

In 1864, German chemist Julius Lothar Meyer, published a table with 28 elements arranged by atomic weight but they did not exactly fit the observed periodicity in chemical properties so he prioritized valency over minor differences in atomic weight.

Between 1863 and 1866, English chemist John Newlands noticed that a list of elements in order of increasing atomic weight, showed a recurrence of similar physical and chemical properties at intervals of eight which coincidentally resembled the octaves of music. Although this idea was ridiculed, he was able to predict the existence of missing elements, such as germanium.

In 1869 and 1870, Dmitri Mendeleev, a Russian chemistry professor, and Julius Lothar Meyer, a German chemist, both published periodic tables with the 64 elements known at that time. They were grouped by atomic weight and properties but Mendeleev occasionally moved elements, such as tellurium and iodine, to better fit with chemical families which produced gaps in the table for element that appeared to be missing. This allowed Mendeleev to accurately predict the properties of several missing elements such as gallium and germanium.

After Ernest Rutherford discovered the atomic nucleus in 1911, Henry Moseley, used the recently invented X-ray spectroscope to confirm that the number of protons in an atom was identical to the place of each element in the periodic table. In 1913 Moseley predicted four of the missing elements:- 43 (technetium), 61 (promethium), 72 (hafnium), and 75 (rhenium), all of which were later discovered. (Modern quantum theories of electron configurations within atoms, made it clear that each row in the periodic table corresponded to the filling of a quantum (valence) shell of electrons).