You can think of an atom's electron affinity as a measure of the attraction that exists between the nucleus, which is positively charged, and the electron, which is negatively charged. This implies that factors that tend to reduce this attraction will also reduce electron affinity. An increase in atomic size leads to a decrease in electron affinity because the incoming electron is added further away from the nucleus, i.
As you go down a group, the outermost electrons are located further and further away from the nucleus. This implies that they feel less pull coming from the nucleus.
Screening plays an important role here, too. On the other hand, an increase in effective nuclear charge has the exact opposite effect. The effective nuclear charge is a measure of the pull exerted on the electrons by the nucleus. In essence, it relates to an increase in atomic number , i. When you add more protons to the nucleus, it will attract electrons more.
If those electrons are added at the same distance from this increasingly powerful nucleus, then they will "feel" more and more pull from it. When you go across a period of the periodic table , atomic number increases, but electrons are being added to the same energy level.
It should reasonably become LESS favourable as we descend a Group of the Periodic Table inasmuch as the nuclear charge becomes more shielded by the intervening electrons, and nuclear attraction to the valence electron should decrease. Thus electron affinity should increase across a Period, but decrease down a Group. Thus we can rationalize the trend on the basis of simple electrostatics. As a chemist, however, you should consult the data in the table provided and elsewhere , and satisfy yourself that I am not telling pork pies!
So why are the halogens the atoms with the greatest electron affinity? How would you explain why the trend of electron affinity is increasing from left to right and decreasing from up to down? Nov 26, Related questions What units is electron affinity measured in? Why does nitrogen have no electron affinity? Nitrogen has a half-filled 2p subshell, so that there is one electron in each orbital. This creates an unusually stable atom because of half-shell stability. Because nitrogen is relatively stable on its own, it has a relatively low electron affinity.
What is the difference between electron affinity and electronegativity? Electronegativity is a chemical property that says how well an atom can attract electrons towards itself. The electron affinity of an atom or molecule is defined as the amount of energy released when an electron is added to a neutral atom or molecule in the gaseous state to form a negative ion.
What is the trend in ionization energy down a group? Moving left to right across a period, atomic radius decreases, so electrons are more attracted to the closer nucleus. The general trend is for ionization energy to decrease moving from top to bottom down a periodic table group.
Moving down a group, a valence shell is added. What happens to ionization energy across a group? Moving left to right within a period or upward within a group, the first ionization energy generally increases. Conversely, 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 greater shielding.
Why does ionization energy increase up a group? The ionization energy of the elements increases as one moves up a given group because the electrons are held in lower-energy orbitals, closer to the nucleus and thus more tightly bound harder to remove.
Which element has the highest ionization energy? From this trend, Cesium is said to have the lowest ionization energy and Fluorine is said to have the highest ionization energy with the exception of Helium and Neon.
Why does reactivity increase down a group? These react by losing electrons and reactivity increases as you go down the group. This is because the increased number of electron shells results in more shielding and a greater distance between the outer electrons and the nucleus, which reduces the attraction of the electrons to the nucleus.
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