When is electron affinity positive or negative

Generally, nonmetals have more positive E ea than metals. Atoms, such as Group 7 elements, whose anions are more stable than neutral atoms have a higher E ea. The electron affinities of the noble gases have not been conclusively measured, so they may or may not have slightly negative values.

Chlorine has the highest E ea while mercury has the lowest. E ea generally increases across a period row in the periodic table, due to the filling of the valence shell of the atom. For instance, within the same period, a Group atom releases more energy than a Group-1 atom upon gaining an electron because the added electron creates a filled valence shell and therefore is more stable.

A trend of decreasing E ea down the groups in the periodic table would be expected, since the additional electron is entering an orbital farther away from the nucleus. Since this electron is farther away, it should be less attracted to the nucleus and release less energy when added. However, this trend applies only to Group-1 atoms.

Electron affinity follows the trend of electronegativity: fluorine F has a higher electron affinity than oxygen O , and so on. The trends noted here are very similar to those in ionization energy and change for similar though opposing reasons. Boundless vets and curates high-quality, openly licensed content from around the Internet. First electron affinities have negative values. For example, the first electron affinity of chlorine is kJ mol By convention, the negative sign shows a release of energy.

When an electron is added to a metal element, energy is needed to gain that electron endothermic reaction. Metals have a less likely chance to gain electrons because it is easier to lose their valance electrons and form cations. It is easier to lose their valence electrons because metals' nuclei do not have a strong pull on their valence electrons. Thus, metals are known to have lower electron affinities.

This trend of lower electron affinities for metals is described by the Group 1 metals:. When nonmetals gain electrons, the energy change is usually negative because they give off energy to form an anion exothermic process ; thus, the electron affinity will be negative. Nonmetals have a greater electron affinity than metals because of their atomic structures: first, nonmetals have more valence electrons than metals do, thus it is easier for the nonmetals to gain electrons to fulfill a stable octet and secondly, the valence electron shell is closer to the nucleus, thus it is harder to remove an electron and it easier to attract electrons from other elements especially metals.

Thus, nonmetals have a higher electron affinity than metals, meaning they are more likely to gain electrons than atoms with a lower electron affinity. For example, nonmetals like the elements in the halogens series in Group 17 have a higher electron affinity than the metals. This trend is described as below. Notice the negative sign for the electron affinity which shows that energy is released.

As the name suggests, electron affinity is the ability of an atom to accept an electron. Unlike electronegativity, electron affinity is a quantitative measurement of the energy change that occurs when an electron is added to a neutral gas atom. The more negative the electron affinity value, the higher an atom's affinity for electrons. 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. A fluorine atom has an electronic structure of 1s 2 2s 2 2px 2 2py 2 2pz 1.

It has 9 protons in the nucleus. The incoming electron enters the 2-level, and is screened from the nucleus by the two 1s 2 electrons. In contrast, chlorine has the electronic structure 1s 2 2s 2 2p 6 3s 2 3p x 2 3p y 2 3p z 1 with 17 protons in the nucleus. There is also a small amount of screening by the 2s electrons in fluorine and by the 3s electrons in chlorine.

This will be approximately the same in both these cases and so does not affect the argument in any way apart from complicating it! Cl Ar K Ca Sc xx. Ti xx. Cr xx. Mn xx. Fe xx. Co xx. Ni xx.

Cu xx. Zn xx. Ga Ge As Se Br Kr Rb Sr Zr xx. Nb xx. Mo xx. Tc xx. Ru xx. Rh xx. Pd xx. Ag xx. Cd xx. In Sn Sb Te I Xe Cs Ba La xx.