Nuclear

Part 7 - Periodic Tables

In 1794, the French chemist Joseph Proust published a paper claiming that elements always reacted with other elements in certain proportions (Proust's Law). He combined copper, carbon and oxygen to make copper carbonate and proved this required the same proportion of weights of the three elements. Likewise, one part oxygen always reacted with one part of mercury to form mercury calx (oxide) and one part oxygen always reacted with two parts of hydrogen to form water. Some substances could combine in new compounds while others could be separated to form simpler ones, while 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? In a 1808 paper, Dalton published the first modern atomic theory. In it he proposed that all matter was composed of indivisible and tiny, indestructible atoms. 

These were in constant states of motion, so tightly packed in solids that they could only vibrate but free to move around in liquids at a speed proportional to the temperature. In the gas state, they were even less densely packed so that they floated freely around and could mix with other gases, as Daniel Bernoulli had proposed in 1738. They were free to move in any direction randomly. 

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. When a solid (like ice) was heated it was typically forced into the liquid state (water) and when heated further it was forced into a gaseous state (water vapour or steam). And this process was reversible as the steam was cooled it condensed into water and when cooled further it solidified into ice.

Every atom of an element was identical to every other atom of the same element. Atoms of different elements, such as oxygen and mercury, had different properties and could be identified by 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. Atoms were unchangeable so chemical reactions did not destroy or alter atoms. Compounds, such as water and mercury calx (mercury oxide), were formed when one atom chemically combined with one or more atoms of other elements.

Elements formed 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).However, Dalton mistakenly believed that the simplest molecule (compound) containing any two elements was always one atom of each (so he thought water was HO, not H2O). Although, by1806, other scientists had already concluded that water molecules contained one oxygen and two hydrogen atoms and that the oxygen atom must weigh 16 times more than hydrogen.  He also did not realize that some atoms typically existed as molecules. (For example, pure free oxygen exists only as a molecule containing two oxygen atoms (O2)). 

An Italian scientist, Amedeo Avogadro corrected the flaws in Dalton's theory in 1811. He proposed that, at equal temperature and pressure, equal volumes of any two gases contained an equal number of molecules. From the volumes at which they reacted, he was able to show the diatomic form (existing as molecules containing two atoms) of many gases and was, thus, able to more accurately estimate the atomic weight of oxygen and other elements, and made a clear distinction between atoms and combination of atoms (molecules).

In 1829, Johann Wolfgang Döbereiner noticed that lithium, sodium and potassium were all soft, reactive metals and deduced that many elements could be grouped together based on their chemical properties.

In 1857, German chemist August Kekulé noted 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 that when he arranged the known elements in a helix on a cylinder by order of increasing atomic weight, he found that elements with similar properties seemed to occur at regular intervals.

In 1864, German chemist Julius Lothar Meyer arranged 28 elements by atomic weight but he prioritized valency over small differences in atomic weight because the elements did not exactly fit the observed similarities in chemical properties.

Between 1863 and 1866, English chemist John Newlands listed elements in order of increasing atomic weight and found that similar physical and chemical properties recurred at intervals of eight, curiously resembling musical octaves. Although the idea was ridiculed, he correctly predicted missing elements, such as germanium.

In 1869 and 1870, Dmitri Mendeleev, a Russian chemistry professor, and Julius Lothar Meyer, a German chemist, independently published periodic tables. They both grouped the 64 elements, known at that time, by atomic weight and properties. However, Mendeleev moved elements, such as tellurium and iodine, as they were more akin to other chemical families. This produced gaps in the table and some element appeared to be missing, so Mendeleev was able to accurately predict the properties of the missing elements 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. This ground breaking work should have earned him a Nobel prize in 1916, but sadly he was killed in 1915, during WW-1. He was just 27.

(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).