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If the first five ionization energies of an element are, respectively: $\pu{1.09 kJ/mol}$, $\pu{2.35 kJ/mol}$, $\pu{4.62 kJ/mol}$, $\pu{6.22 kJ/mol}$ and $\pu{37.83 kJ/mol}$, to which group of the periodic table does this element belong? Predict the valence configuration of this element. Justify your reasoning

I don't understand how to solve this. How do I relate ionization energy to finding the element, then predicting the valence configuration?

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    $\begingroup$ Those ionization values are too small. Did you copy them correctly? $\endgroup$
    – Buck Thorn
    Mar 1, 2021 at 8:07

1 Answer 1

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You expect ionization energies to increase as you remove more electrons from an atom, because the charge of the nucleus remains the same (and therefore its attraction to the remaining electrons) but repulsion from other electrons is reduced with removal of each additional electron. However, in addition to this simple intuitive trend, quantum mechanics throws in a twist, namely the existence of electron "shells" about an atom, sharing similar energy (the same principal quantum number). Electrons in the same shell repel each other with the same intensity. This is the justification for Slater's Rules.

There is a third way to solve this: look up values in a table of ionization energies. It helps ensure accuracy to have "second and third opinions", so this is a good exercise. From the Wikipedia (in $\pu{kJ/mol}$), the values are 1086.5, 2352.6, 4620.5, 6222.7 37831, 47277.0.

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