Following Dobereiner, in 1864 John Newlands, an English scientist found that placed in order of increasing atomic mass, elements which has similar physical and chemical properties occurred at every eight interval. This was called Newlands’ Law of Octaves, however this came under scrutiny as elements above calcium did not fit the pattern, and as more elements such as the Nobel gasses were discovered, there was no place for them in the table.
In 1869 Dimitrii Mendeleev, a Russian and four months later in 1870, Lothar Meyer, a German, independently produced the periodic table. Both Mendeleev and Meyer ordered the elements by increasing atomic mass and both left gaps where they predicted elements were missing. Mendeleev’s periodic table gave a chart of elements that were grouped so elements showing similar properties occurred in the same vertical group. This included similar melting and boiling points, and a similar valency. Mendeleev (1869 pp-32-33) states that “the elements, if arranged according to their atomic weights, exhibit an evident periodicity of properties’’, hence the name, the periodic table.
Mendeleev also, stated that if the atomic weight of an element caused it to be placed in the wrong group then the weight of that element must be wrong. He therefore corrected the atomic masses of beryllium (Be), indium (In) and uranium (U). He was so confident with the accuracy of his table that he used it to predict the physical properties of three elements that at the time were unknown. Between 1874 and 1885, the elements which Mendeleev had predicted, scandium, gallium and germanium were discovered and were found to be extremely close to the values given. Because of this Mendeleev’s table was generally accepted and he became known as the ‘Father of modern chemistry’.
Mendeleev’s periodic table was very successful, although did show to have a few problems; Firstly, the positions of isotopes could not be placed in the table. As an example, normal carbon or 12C would fit perfectly in the table. An isotope of carbon, 14C however, exhibits the same characteristics as 12C would not fit in the table as it would have to be accommodated along with nitrogen, in the same position. Secondly, in order for the elements to fit the requirements of the Mendeleevillian table, those in a particular column have the same valance. Mendeleev in certain cases put an element which had a higher atomic weight ahead of one which had a slightly lower atomic weight. To keep tellurium (Te) in the valance 2 column and iodine (I) in the valance 1 column, tellurium was put ahead of iodine, even though it had a higher atomic weight. This was to later make part of the modern periodic table.
The successes of the table were that although when Mendeleev presented his table, the Nobel gases were still undiscovered, as they were found, they could be put neatly as the last group of elements without disturbing the rest of the table. The major success was due to his arrangement of the table, predicting elements which were missing was an accurate process. Even now, with 112 elements found (both naturally occurring and artificially made) we are still able to predict the properties of elements by looking vertically at groups and horizontally across the periods.
On the foundations laid down by Mendeleev, in 1913, Henry Mosley, using his work with x-rays was able to determine the precise atomic number of the elements. Thus, he rearranged the table in order of increasing atomic number. The vertical groups in the table all have similar outer electron arrangements which determine the number of chemical bonds which an atom of that element can make. Elements which are seen horizontally in the periods show an increase in valence the further to the right of the table they are. For example, group 1, the alkali metals have a single valence electron and display similar properties of chemical reactivity such as the formation of oxides, and their reaction with water. The elements in a period show increase in the last electron configuration.
In all, the modern periodic table comprises of 7 periods which increase in length as the order of the period increases. The elements in a period have consecutive atomic numbers. The first period is the shortest, containing just two elements, hydrogen (H) and helium (He). The second and third periods each contain eight elements, the fourth and fifth contain eighteen elements while the sixth period contains thirty-two elements. Fifteen of these elements (57-71) are separated from the table to form the lanthanide series of rare earth metals which display similar properties. The seventh and final period contains the rest of the elements and is as of yet, incomplete. Again, this period has a series of fifteen elements which are separated from the table, called the actinide series. The vertical columns or groups total 18 which are divided into A and B groups. Groups 1A to VIIIA hold the main group elements while 1B to VIIIB holds the transition metals. The lanthanide and actinide series are known as the inner-transition elements.
The modern periodic table is roughly split into metals and non-metals. Those which exhibit the most metallic properties such as the alkali metals are placed on the far left of the table while the non-metals occupy the right of the table. The Nobel gases which are completely inert due to having full outer electron shells are set at the far right of the table in group 18. The transition elements, which are a bridge between highly metallic alkali elements and the non-metals, lie in the centre of the table and exhibit characteristics of both metals and non-metals.
As the elements move from left to right across a period, the metallic character or the tendency to lose electrons decreases but as they move down a group, this tends to increase. This is due to the fact that there is a general increase in the ionisation energy moving across a period whereas there is a decrease moving down the group. As we move from left to right, we see the reverse and the elements tend to become more metallic, losing their non-metallic properties. Similarly, as they go down a group, non metallic character tends to decrease.
References
On the Relationship of the Properties of the Elements to their Atomic Weights D. Mendeleev,
Fundamentals of Chemistry, Second Edition, David E. Goldberg (1998)
Website: http://home.att.net/~cat6a/class_elem-V.htm
Course Notes
