Topic 7 - Periodicity
Dobereiner - Law of Triads
Early in the 19th century, scientists began to seek ways to classify the elements. One attempt was by Johann Dobereiner, a German chemist, in 1817. Dobereiner found that properties of calcium, barium, and strontium were very similar. He also noted that the atomic mass of strontium was about midway between those of calcium and barium. He grouped these three elements into what he termed a triad . Later, Dobereiner found several other groups of three elements with similar properties. In each case the middle element has an atomic mass about halfway between those of the first and third elements in the triad. Dobereiner's ideas were not taken seriously by others - they looked like coincidences, but the idea of trends had begun.
Newlands - Law of Octaves
In 1863 John Newlands, an English chemist, suggested a classification. He arranged the elements in order of their increasing atomic masses. He noted there appeared to be a repetition of similar properties every eighth element. Therefore, he placed seven elements in each group. He then arranged the 49 elements known at that time into seven groups of seven each. Newlands referred to his arrangement as the Law of Octaves because the same properties repeated every eight elements. Newlands' arrangement worked very well through the element Ca, after that the properties did not seem to match up as well with the vertical groups. For example, in the O, S, Fe, Se group, iron (Fe) does not belong in the same group as oxygen and sulfur. This left room for improvement. The obviously misplaced elements are in bold.
Mendeleev
In 1869, Dmitri Mendeleev, a Russian chemist, proposed a similar idea. He suggested that the properties of the elements were a function of their atomic masses. However, Mendeleev believed that similar properties occurred after periods (horizontal rows) that could vary in length. Although he placed seven elements in each of his first two rows, he placed seventeen elements in each of the next two. However Mendeleev took it one step further in predicting that some elements may not be discovered yet and left blank spots in the periodic table. He even went so far as to predict a name and the properties of these yet to be discovered elements. It later turned out he was very close on his predictions. Below is a version of his early periodic table. Mendeleev stated that the properties of the elements are a periodic function of their atomic masses. This statement is called the periodic law. There was a problem with Mendeleev's table of elements - tellurium and iodine seemed in the wrong columns if they were placed only by atomic mass. However Mendeleev assumed the mass was measured incorrectly and placed them according to their properties. He figured that later measurements would prove him correct.
Moseley
More inconsistencies were found with early versions of the periodic table. Henry Moseley, in 1913, found the reason for the inconsistencies. He performed X-ray experiments on the elements and found that each element has an integral positive charge, the atomic number. As a result of Moseley's work, the periodic law was revised. It now has as its basis the atomic numbers of the elements instead of the atomic masses. The modern statement of the periodic law is the properties of the elements are a periodic function of their atomic numbers. The atomic number indicates the number of protons in the nucleus of each atom of the element. Because an atom is electrically neutral, the atomic number also indicates the number of electrons surrounding the nucleus. Certain electron arrangements are repeated periodically as atoms increase in atomic number. As you have already seen, atoms with similar electron configurations are placed in the same column. This table is called the periodic table of the elements.
Patterns of first ionisation energies in the Periodic Table
First ionisation energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar. These variations in first ionisation energy can all be explained in terms of the structures of the atoms involved.
Factors affecting the size of ionisation energy
Ionisation energy is a measure of the energy needed to pull a particular electron away from the attraction of the nucleus. A high value of ionisation energy shows a high attraction between the electron and the nucleus. The size of that attraction will be governed by:
The charge on the nucleus - The more protons there are in the nucleus, the more positively charged the nucleus is, and the more strongly electrons are attracted to it.
The distance of the electron from the nucleus - Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.
The number of electrons between the outer electrons and the nucleus - Consider a sodium atom, with the electronic structure 2,8,1. (There's no reason why you can't use this notation if it's useful!) If the outer electron looks in towards the nucleus, it doesn't see the nucleus sharply. Between it and the nucleus there are the two layers of electrons in the first and second levels. The 11 protons in the sodium's nucleus have their effect cut down by the 10 inner electrons. The outer electron therefore only feels a net pull of approximately 1+ from the centre. This lessening of the pull of the nucleus by inner electrons is known as screening or shielding.
Measures of atomic radius
Unlike a ball, an atom doesn't have a fixed radius. The radius of an atom can only be found by measuring the distance between the nuclei of two touching atoms, and then halving that distance.
The same atom could be found to have a different radius depending on what was around it. The left hand diagram shows bonded atoms. The atoms are pulled closely together and so the measured radius is less than if they are just touching. This is what you would get if you had metal atoms in a metallic structure, or atoms covalently bonded to each other. The type of atomic radius being measured here is called the metallic radius or the covalent radius depending on the bonding. The right hand diagram shows what happens if the atoms are just touching. The attractive forces are much less, and the atoms are essentially "unsquashed". This measure of atomic radius is called the van der Waals radius after the weak attractions present in this situation.
Trends in atomic radius in the Periodic Table
The exact pattern you get depends on which measure of atomic radius you use - but the trends are still valid. The following diagram uses metallic radii for metallic elements, covalent radii for elements that form covalent bonds, and van der Waals radii for those (like the noble gases) which don't form bonds.
Trends in atomic radius down a group
It is fairly obvious that the atoms get bigger as you go down groups. The reason is equally obvious - you are adding extra layers of electrons.
Trends in atomic radius across periods
You have to ignore the noble gas at the end of each period. Because neon and argon don't form bonds, you can only measure their van der Waals radius - a case where the atom is pretty well "unsquashed". All the other atoms are being measured where their atomic radius is being lessened by strong attractions. You aren't comparing like with like if you include the noble gases.
If you think about it, the metallic or covalent radius is going to be a measure of the distance from the nucleus to the electrons which make up the bond. From lithium to fluorine, those electrons are all in the 2-level, being screened by the 1s2 electrons. The increasing number of protons in the nucleus as you go across the period pulls the electrons in more tightly. The amount of screening is constant for all of these elements. In the period from sodium to chlorine, the same thing happens. The size of the atom is controlled by the 3-level bonding electrons being pulled closer to the nucleus by increasing numbers of protons - in each case, screened by the 1- and 2-level electrons.
