The periodic table organizes the chemical elements according to periodic trends in the chemical activity which has been discovered over centuries of research and discovery. The main value of the periodic table in chemistry is that the chemical properties of an element can be predicted based on its location on the table.
The first recognizable version of the periodic table was prepared by Dimitri Mendeleev in 1869. He listed in a row the elements in order according to their atomic weight, and started a new row under the first when the next element had similar chemical properties to the first member of the row. For example, hydrogen, lithium, sodium, and potassium have similar chemical properties, so they are placed in the same column.
This simple approach toward organizing the chemical properties of the elements produced some puzzling groupings between elements that did not seem to belong together. For example, this approach would group silicon and titanium, which do not seem to share any common properties. However, the trend toward periodic behavior was so strong that Mendeleev made some changes in the grouping aimed at making the periodic behavior more apparent.
The first was to leave gaps in the table when none of the known elements matched. Remember that discovery of new elements was rapidly progressing in the middle of the 19th Century, so that the idea that there might be unknown elements to add to the mix was quite reasonable. Mendeleev was sufficiently confident in the essential periodicity of the elements that he predicted the chemical properties of the missing elements – correctly, as it turned out.
Mendeleev also recognized that there were some atoms which, when ordered by atomic weights, seemed to be reversed; that is, their chemical properties suggested that they should be in each other’s group. He switched these elements, abandoning the strict ordering by atomic weight in favor of his faith in chemical periodicity. As more became known about atomic structure, scientists realized that by following the guide of chemical periodicity, Mendeleev had placed the elements in order of the number of protons in the nucleus. This was a particularly impressive achievement in that neither the nucleus nor protons were known to exist at that time.
Let’s take a closer look at some of these periodicities which are displayed in the periodic table. Elements in the periodic table appear in columns called groups and rows called periods. In the standard periodic table there are 18 groups and 7 periods, with the f-shell lanthanides and actinides being listed separately.
Elements in the same group exhibit trends in their atomic radium, ionization energy, and electronegativity (ability of an atom to accept an additional electron). The atomic radius increases as the atomic number of a group element increases. Since the outer electrons are situated further from the nucleus, it is easier to remove an electron, giving lower ionization energy in larger group elements. A similar trend appears in electronegativity, which decreases in larger atoms owing to the larger separation between the valence electrons and the nucleus.
Elements from the same period (a row from left to right) exhibit trends in these same chemical properties. Period elements have decreasing atomic radius as their atomic number increases. The smaller atomic radii then give an increase in ionization energy across the period. Electronegativity also increases in a period from left to right because the electrons reside closer to the nucleus.
Periodic tables today often are annotated with these additional chemical properties, which helps one improve the accuracy of chemical reactivity predictions. However, one can make surprisingly accurate predictions based only on relative positions of species in the periodic table.
For example, say you need to know the melting point of sodium fluoride, and you know that the melting point of sodium chloride is 801 degrees Centigrade. Well, the periodic table tells us that fluorine has a larger electronegativity than chlorine, which is positioned below fluorine in the same group. The large difference in electronegativity between sodium and the halides tells us that their bonding is ionic, and the strength of an ionic bond increases as the difference in electronegativity increases. We therefore predict that the sodium fluoride is more strongly bound than is sodium chloride, thus the melting point of sodium fluoride will be somewhat larger than that of sodium chloride. Indeed, sodium fluoride melts at 993 degrees Centigrade, so the prediction is correct.
The periodic table provides a remarkably concise summary of chemical property trends among the elements. Still more impressive is that the organizing principles developed by Mendeleev produced order in chemical properties unknown at that time and not understood until 80 years later. This work of genius revolutionizes the practice of chemistry worldwide.
