The properties of the elements exhibit trends and these trends can be predicted with the help of the periodic table. They can also be explained and understood by analyzing the electron configurations of the elements. This is because, elements tend to gain or lose valence electrons to achieve the stable octet formation.
The properties of the elements exhibit trends and these trends can be predicted with the help of the periodic table. They can also be explained and understood by analyzing the electron configurations of the elements. This is because, elements tend to gain or lose valence electrons to achieve the stable octet formation.
In addition to this activity, there are two other important trends. First, electrons are added, one at a time, moving from left to right across a period. And, as this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction. As a result, the electrons become closer to the nucleus and more tightly bound. The second trend is the moving down a column in the periodic table, where the outermost electrons become less tightly bound to the nucleus. And these trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegativity.
But, before going into that we need to know a bit more about the above mentioned terms:
Atomic Radius
The atomic radius of an element is half of the distance between the centers of two atoms of an element that are in contact with each other. Generally, the atomic radius decreases across a period, from left to right and increases down a given group. Therefore, the atoms with the largest atomic radii are located in Group I and at the lower half of groups.
Ionization Energy
Ionization energy or ionization potential is the energy required to completely remove an electron from a gaseous atom or ion. And, the closer and more tightly an electron is bound to the nucleus, the more difficult it is to remove and the higher its ionization energy. Ionization energy is also required to remove a second valence electron from the univalent ion to form the divalent ion, and so on.
Electron Affinity
Electron affinity is the energy change that occurs when an electron is added to a gaseous atom. It reflects the ability of an atom to accept an electron. And the atoms with stronger effective nuclear charge have a greater electron affinity. Therefore, some generalizations can be made about the electron affinities of certain groups in the periodic table. The alkaline earths have low electron affinity values. This is because they have filled sub shells. But, the halogens have high electron affinities because of the addition of an electron to an atom results in a completely filled shell. Noble gases have zero electron affinities, since each atom possesses a stable octet and will not accept an electron readily.
Electro negativity
An atom with higher electro negativity has a great capacity for attracting bonding electrons. Therefore, electro negativity is a measure of the attraction of an atom for the electrons in a chemical bond. It’s related to ionization energy. So, electrons with low ionization energies have low electro negativities because their nuclei do not exert a strong attractive force on electrons. And, elements with high ionization energies have high electro negativities. This is because of the strong pull exerted on electrons by the nucleus.
Therefore, electro negativity is dependant on the atomic number. As the atomic number increases, the electro negativity decreases, as a result of increased distance between the valence electron and nucleus. An example of an electropositive element, i.e. one with low electro negativity, is cesium. And an example of a highly electronegative element is fluorine.
About the author:
Dr.George Grant is an experienced researcher in Bio-chemistry. He has done extensive researches and experiments in the field. He is a visiting faculty for some of the most reputed Science colleges. For more information on Chemistry Tools and Definitions, Please Visit- http://www.chemicool.com
