There is a gradual increase in the 1 st e i from left

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There is a gradual increase in the 1 st E i from left to right across a row of the periodic table. The group 18 elements (noble elements) have filled valence orbitals , and a relatively high Z eff because electrons within the same subshell don’t shield one another very strongly, hence large 1 st E i . Group 1 elements , have a single valence electron that is well shielded from the nucleus by all the inner-shell electrons, called the core electrons , resulting in a low Z eff . The valence electron is thus held loosely, hence small 1 st E i .
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E i gradually decrease going down a group in the periodic table, e.g. from Li to Fr. o As atomic number increases going down a group, both the principal quantum number of the valence-shell electrons and their average distance from the nucleus increases. o As a result, the valence-shell electrons are less tightly held, and E i decreases. What about the anomalies to these general trends?? Going across the period E i of Be > B and N > O. Likewise Mg > Al and P > S. Why is that??
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Anomalies in 1 st E i going from group 2 to group 3 A 2 s electron is removed on ionization of Be, but a 2 p electron is removed on ionization of B. A 2 s electron is on average closer to the nucleus than a 2 p electron , it is held more tightly and is harder to remove . Thus, E i of Be > B. Or equivalently the 2 p electron of B is shielded somewhat by the 2 s electrons , feels a smaller Z eff and is thus more easily removed than a 2 s electron of Be.
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Anomalies in 1 st E i going from group 15 to group 16 In comparing N with O, for example, the N electron is removed from a half-filled orbital , vs a filled orbital in O. Because electrons repel one another and tend to stay as far apart as possible, electrons that are forced together in a filled orbital are slightly higher in energy than those in a half-filled orbital , so it is slightly easier to remove one . Thus, O has a smaller E i than N.
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Successive ionizations Successively larger amounts of energy are required for each successive ionization step because it is much harder to remove a negatively charged electron from a positively charged ion than from a neutral atom. Give a short explanation for the extremely large values observed in the pink regions of the table.
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Electron Affinity and Electronegativity Just as an electron can be removed form an atom to form a cation, it’s also possible add an electron to an atom to form an anion. An element’s electron affinity , abbreviated E ea , is the energy change that occurs when an electron is added to an isolated atom in the gaseous state . Ionization energies are always positive because energy must always be added to remove an electron from an atom. Electron affinities , however, are generally negative because energy is usually released when a neutral atom adds an electron.
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