Mandeleev (1834-1907) is accredited to have successfully discovered laws governing the periodic nature of atoms and elements in the 1860s. A chemical element refers to the simplest chemical substance that cannot be broken down chemically with a distinct number of protons (atomic number) whereas an atom refers to the smallest particle of an element which has the same basic chemical properties to its element. Mendeleev discovered that similar chemical properties were exhibited by elements when arranged in order of their increasing atomic weight (nowadays periods and groups) and he termed this trend as periodicity (Spronsen, 1969).
Properties of elements in relation to their atomic weights
Rowsperiods of the periodic table consist of elements with the same quantum number in the outermost orbital whereas columnsgroups consist of elements with the same electronic configuration in the outermost orbital. Across the period, the force exerted upon the outermost electrons increases with an increase of atomic number because there are more protons in the nucleus. In contrast, downward the group, the force decreases because of the shielding effect exerted by lower energy levels on the nucleus. Hence, the atomic radius increases down the group compared to a decrease across the period. For instance, lithium and fluorine are in the same period whereby Lithium has three protons compared to nine protons for Fluorine. Because of higher nuclear charge in Fluorine than in Lithium, the outermost electronsshells are strongly bound to the nucleus than in Lithium which renders Lithium atom to have a large atomic radius than the former even though it has fewer protons. In another other case, Fluorine and Chlorine are in the same group but the Chlorine has a larger atomic radius due to an increased number of protons (and shells) than Flourine. These two trends of atomic radius across and down the group determine the properties of elements to a large extent (Wetzel, 1993).
The second property of atoms and elements is the ionization energy which refers to the amount of energy required to remove an electron from a gaseous atom in its ground state (Ibid, 1993). With a decrease in atomic radius across the period, the ionization energy increases whereas with an increase in atomic radius down the group, the ionization energy decreases. Outermost electrons are held more strongly to the nucleus across the period than down the group. This explains the difference in relative amounts of ionization energy required for elements across the period and down the group. Electronic affinity is the third property which refers to the energy releasedabsorbed when an electron is added to a neutral, gaseous atom (Ibid, 1993). It also depends on the atomic radius down or across the group whereby reactions are more exothermic across the period because the electrons are tightly bound- in contrast to movement down the group. Fourth, electronegativity represents the relative attraction of electrons by the nucleus of a particular atom in a chemical bond and its dependant on electron affinity as well as ionization energy. It increases and decreases across the period and down the group respectively.
Wetzel (1993) in his work sates that the outermost electronic shells of noble gases (Helium, Neon etc) are completely occupied hence they cannot attract andor lose electron(s) are chemically inert. On the other hand, alkali metals (Lithium, Sodium etc) consist of one valence electron in the outermost shells whereas halogens have seven valence electrons. Therefore alkali metals have high affinity, are easily ionized, and highly reactive because they lose only one electron to acquire a stable state. The same case applies to halogens (Fluorine, Chlorine etc) which need only one electron to fill up the outermost shell and acquire a stable state. Ionization energy of halogens is therefore high with respect to their relatively small atomic radius besides their high affinity.
Properties of elements in relation to their atomic weights
Rowsperiods of the periodic table consist of elements with the same quantum number in the outermost orbital whereas columnsgroups consist of elements with the same electronic configuration in the outermost orbital. Across the period, the force exerted upon the outermost electrons increases with an increase of atomic number because there are more protons in the nucleus. In contrast, downward the group, the force decreases because of the shielding effect exerted by lower energy levels on the nucleus. Hence, the atomic radius increases down the group compared to a decrease across the period. For instance, lithium and fluorine are in the same period whereby Lithium has three protons compared to nine protons for Fluorine. Because of higher nuclear charge in Fluorine than in Lithium, the outermost electronsshells are strongly bound to the nucleus than in Lithium which renders Lithium atom to have a large atomic radius than the former even though it has fewer protons. In another other case, Fluorine and Chlorine are in the same group but the Chlorine has a larger atomic radius due to an increased number of protons (and shells) than Flourine. These two trends of atomic radius across and down the group determine the properties of elements to a large extent (Wetzel, 1993).
The second property of atoms and elements is the ionization energy which refers to the amount of energy required to remove an electron from a gaseous atom in its ground state (Ibid, 1993). With a decrease in atomic radius across the period, the ionization energy increases whereas with an increase in atomic radius down the group, the ionization energy decreases. Outermost electrons are held more strongly to the nucleus across the period than down the group. This explains the difference in relative amounts of ionization energy required for elements across the period and down the group. Electronic affinity is the third property which refers to the energy releasedabsorbed when an electron is added to a neutral, gaseous atom (Ibid, 1993). It also depends on the atomic radius down or across the group whereby reactions are more exothermic across the period because the electrons are tightly bound- in contrast to movement down the group. Fourth, electronegativity represents the relative attraction of electrons by the nucleus of a particular atom in a chemical bond and its dependant on electron affinity as well as ionization energy. It increases and decreases across the period and down the group respectively.
Wetzel (1993) in his work sates that the outermost electronic shells of noble gases (Helium, Neon etc) are completely occupied hence they cannot attract andor lose electron(s) are chemically inert. On the other hand, alkali metals (Lithium, Sodium etc) consist of one valence electron in the outermost shells whereas halogens have seven valence electrons. Therefore alkali metals have high affinity, are easily ionized, and highly reactive because they lose only one electron to acquire a stable state. The same case applies to halogens (Fluorine, Chlorine etc) which need only one electron to fill up the outermost shell and acquire a stable state. Ionization energy of halogens is therefore high with respect to their relatively small atomic radius besides their high affinity.
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