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Chapter 4: Problem 2

For atoms with many electrons, not all electrons are at the same distance from the nucleus. In this case, which electron would have the lowest ionization energy: the electron that is closest to the nucleus or the electron that is farthest from the nucleus? Explain.

Short Answer

Expert verified
The electron that is farthest from the nucleus would have the lowest ionization energy, because there is less attraction between it and the nucleus.

Step by step solution

01

Understanding Ionization Energy

Ionization energy refers to the minimum amount of energy needed to remove an electron from an atom or ion. This energy is directly proportional to the strength of the attraction between the electron and the nucleus: the stronger the attraction, the more ionization energy is needed to remove the electron.
02

Identifying Electron Positions

Electron positions in the atom are defined by their energy levels (or shells). Electrons closer to the nucleus occupy lower energy levels and are more strongly attracted to the positively charged protons in the nucleus. On the other hand, electrons farther from the nucleus are on higher energy levels and face less attraction force from the nucleus.
03

Determining Ionization Energies

Given that ionization energy is directly proportional to the strength of the attraction between the electron and nucleus, electrons closer to the nucleus would require more energy to remove due to the stronger attraction force. Therefore, the electron closest to the nucleus has the highest ionization energy. Conversely, the electron farthest from the nucleus faces less attraction and would require less energy to remove, hence it has the lowest ionization energy.

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Most popular questions from this chapter

a) How much energy does it take to remove an electron from one \(\mathrm{H}\) atom? b) How much total energy would it take to remove the electrons from two H atoms? c) How much total energy would it take to remove the electrons from a mole of \(\mathrm{H}\) atoms? Express this energy in units of \(\mathrm{J}\) and in units of \(\mathrm{MJ} .\)

Using grammatically correct sentences: a) provide a possible explanation for why \(\mathrm{IE}_{1}\) for \(\mathrm{He}\) is greater than \(\mathrm{IE}_{1}\) for \(\mathrm{H}\). b) provide a possible explanation for why \(\mathrm{IE}_{1}\) for \(\mathrm{Li}\) is less than \(\mathrm{IE}_{1}\) for He.

Recall that the IE of \(\mathrm{H}\) is \(1.31 \mathrm{MJ} / \mathrm{mole}\). If all three electrons in Li were in the first shell at a distance equal to that of hydrogen, which of the following values would be the better estimate of the IE \(_{1}\) of Li: \(3.6 \mathrm{MJ} / \mathrm{mole}\) or \(0.6 \mathrm{MJ} / \mathrm{mole}\) ? Explain your reasoning.

The value of the ionization energy of He given in Table 1 is described as being consistent with a model in which the two electrons in He are in a "shell" at approximately the same distance from the nucleus as the one electron in \(\mathrm{H}\). Use the Coulombic Potential Energy equation, \(V=\frac{\mathrm{kq}_{1} \mathrm{q}_{2}}{\mathrm{~d}}\) to explain how this conclusion can be reached. Hint: recall the relationship between \(V\) and \(\mathrm{IE}_{1}\).

Based on what you have learned so far in this course, predict the relationship between \(\mathrm{IE}_{1}\) and atomic number by making a rough graph of \(\mathrm{IE}_{1}\) vs. atomic number. Discuss possible ideas with your team and decide which one you think makes the most sense. DO NOT PROCEED TO THE NEXT PAGE UNTIL YOU HAVE MADE YOUR PREDICTED GRAPH.
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