The 14th International Conference on

Miniaturized Systems for Chemistry and Life Sciences

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Periodic trends


Periodic trends

Periodic trends are specific patterns in the properties of chemical elements that are revealed in the periodic table of elements. Major periodic trends include electronegativity, ionization energy, electron affinity, atomic radii, ionic radius, metallic character, and chemical reactivity.

Periodic trends arise from the changes in the atomic structure of the chemical elements within their respective periods (horizontal rows) and groups in the periodic table. These laws enable the chemical elements to be organized in the periodic table based on their atomic structures and properties. Due to the periodic trends the unknown properties of any element can be partially known.

This principle was discovered by Dmitri Mendeleev in 1871 after a number of investigations by scientists in the 19th century. Mendeleev also proposed a periodic system of elements that was based not only on atomic weights but also on chemical and physical properties of the elements and their compounds. In 1913, Henry Moseley determined that periodicity depends on the atomic number rather than atomic weight. Lothar Meyer presented his table several months after Mendeleev, but opposed his Periodic law. Initially, no theoretical explanation for the Periodic Law was available and it was used only as an empirical principle, but, with the development of the quantum mechanics, it became possible to understand the theoretical basis for the Periodic Law.

Discovery of Periodic Law constitutes one of the most singularly important events in the history of chemical science. Almost every chemist makes extensive and continued use of Periodic Law. Periodic Law also led to the development of the periodic table, which is widely used nowadays.

As one progresses down a group on the periodic table, the ionization energy will likely decrease since the valence electrons are farther away from the nucleus and experience a weaker attraction to the nucleus's positive charge. There will be an increase of ionization energy from left to right of a given period and a decrease from top to bottom. As a rule, it requires far less energy to remove an outer-shell electron than an inner-shell electron. As a result, the ionization energies for a given element will increase steadily within a given shell, and when starting on the next shell down will show a drastic jump in ionization energy. Simply put, the lower the principal quantum number, the higher the ionization energy for the electrons within that shell. The exceptions are the elements in the boron and oxygen family, which require slightly less energy than the general trend.

added to it, conversely as the energy required to detach an electron from a singly charged anion. The sign of the electron affinity can be quite confusing, as atoms that become more stable with the addition of an electron (and so are considered to have a higher electron affinity) show a decrease in potential energy; i.e. the energy gained by the atom appears to be negative. For atoms that become less stable upon gaining an electron, potential energy increases, which implies that the atom gains energy. In such a case, the atom's electron affinity value is positive. Consequently, atoms with a more negative electron affinity value are considered to have a higher electron affinity (they are more receptive to gaining electrons), and vice versa. However, in the reverse scenario where electron affinity is defined as the energy required to detach an electron from an anion, the energy value obtained will be of the same magnitude but have the opposite sign. This is because those atoms with a high electron affinity are less inclined to give up an electron, and so take more energy to remove the electron from the atom. In this case, the atom with the more positive energy value has the higher electron affinity. As one progresses from left to right across a period, the electron affinity will increase.

Electronegativity is a measure of the ability of an atom or molecule to attract pairs of electrons in the context of a chemical bond. The type of bond formed is largely determined by the difference in electronegativity between the atoms involved, using the Pauling scale. Trend-wise, as one moves from left to right across a period in the periodic table, the electronegativity increases due to the stronger attraction that the atoms obtain as the nuclear charge increases. Moving down in a group, the electronegativity decreases due to the longer distance between the nucleus and the valence electron shell, thereby decreasing the attraction, making the atom have less of an attraction for electrons or protons.

Valence electrons are the electrons in the outermost electron shell of an isolated atom of an element. Sometimes, it is also regarded as the basis of Modern Periodic Table. In a period, the number of valence electrons increases (mostly for light metal/elements) as we move from left to right side. However, in a group this periodic trend is constant, that is the number of valence electrons remains the same.

The greater the number of core electrons, the greater the shielding of electrons from the core charge of the nucleus. For this reason ionization energy is lower for elements lower down in a group, and polarizability of species is higher for elements lower down in a group. The valency does not change going down a group since the bonding behavior is not affected by the core electrons. However, non-bonding interactions such as those just cited are affected by core electrons.

Metallic properties increase down groups as decreasing attraction between the nuclei and the outermost electrons causes the outermost electrons to be loosely bound and thus able to conduct heat and electricity. Across the period, from left to right, increasing attraction between the nuclei and the outermost electrons causes metallic character to decrease.


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