The correct order of electron affinity is:
(a) F > C1 > Br (b) F> Cl< Br
(c) F< Cl > Br
(d) F< Cl< Br
Answers
Since the atomic size increases down the group, electron affinity generally decreases (At < I < Br < F < Cl). An electron will not be as attracted to the nucleus, resulting in a low electron affinity. However, fluorine has a lower electron affinity than chlorine
Explanation:
Electron affinity is defined as the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion. In other words, the neutral atom's likelihood of gaining an electron.
Periodic Table showing Electron Affinity Trend
Patterns in Electron Affinity
Electron affinity increases upward for the groups and from left to right across periods of a periodic table because the electrons added to energy levels become closer to the nucleus, thus a stronger attraction between the nucleus and its electrons. Remember that greater the distance, the less of an attraction; thus, less energy is released when an electron is added to the outside orbital. In addition, the more valence electrons an element has, the more likely it is to gain electrons to form a stable octet. The less valence electrons an atom has, the least likely it will gain electrons.
Electron affinity decreases down the groups and from right to left across the periods on the periodic table because the electrons are placed in a higher energy level far from the nucleus, thus a decrease from its pull. However, one might think that since the number of valence electrons increase going down the group, the element should be more stable and have higher electron affinity. One fails to account for the shielding affect. As one goes down the period, the shielding effect increases, thus repulsion occurs between the electrons. This is why the attraction between the electron and the nucleus decreases as one goes down the group in the periodic table.
As you go down the group, first electron affinities become less (in the sense that less energy is evolved when the negative ions are formed). Fluorine breaks that pattern, and will have to be accounted for separately. The electron affinity is a measure of the attraction between the incoming electron and the nucleus - the stronger the attraction, the more energy is released. The factors which affect this attraction are exactly the same as those relating to ionization energies - nuclear charge, distance and screening. The increased nuclear charge as you go down the group is offset by extra screening electrons. Each outer electron in effect feels a pull of 7+ from the center of the atom, irrespective of which element you are talking about.
Energy is released when a electron is added to a nonmetal.
Nonmetals have a greater electron affinity than metals because their atomic structure allows them to gain electrons rather than lose them.
Atoms with a low electron affinity want to give up their valence electrons because they are further from the nucleus; as a result, they do not have a strong pull on the valence electrons.
As you move down a group on the periodic table, electron affinity decreases. First, the electrons are placed in energy levels further away from the nucleus, which results in electrons not having a strong attraction to the nucleus; secondly, the atom does not want gain electrons because there is minimal charge on the outer energy levels from the nucleus; and lastly, the shielding effect increases, causing repulsion between the electrons, thus they move further from each other and the nucleus itself.
Nonmetals want to gain electrons because they have more valence electrons than metals, so it is easier for them to gain electrons than lose the valance electrons to fulfill a stable octet. In addition, nonmetals' valance electrons are closer to the nucleus, thus allowing more attraction between the two.
Metals have a low electron affinity (a less likely chance to gain electrons) because they want to give up their valence electrons rather than gain electrons, which require more energy than necessary. In addition, they do not have a strong pull on the valance electrons because they are far away fr
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