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Chapter 9: Problem 83

Write electron configurations to show the first two ionizations for Cs. Explain why the second ionization energy is much greater than the first.

Short Answer

Expert verified
The electron configurations for Cs, Cs+ and Cs++ are [Kr] 5s^1, [Kr], and [Ar] 4d^10 4p^6 respectively. The second ionization energy is higher than the first due to the second electron being removed from a lower energy level (4p), which is closer to and more strongly attracted to the atomic nucleus.

Step by step solution

01

Electron Configuration of Cs

Cesium (Cs) has 55 electrons. It's electron configuration can be written as \([Kr] 5s^1\), where [Kr] represents core electrons identical to the noble gas krypton.
02

First Ionization of Cs

The first ionization of Cs involves removing one electron from the outermost shell. The electron configuration becomes \([Kr]\). The outermost electron is in the 5s orbital, which is relatively easy to remove, thus the lower ionization energy.
03

Second Ionization of Cs

The second ionization of Cs involves removing another electron. However, this time, the electron is being removed from a complete shell, closer to the nucleus (the 4p orbital), represented as [Ar] 4d^10 4p^6. Removing an electron from this shell requires much more energy because these electrons are much closer to the nucleus and thus more strongly attracted to the positive charge of the protons.
04

Reason for Higher Second Ionization Energy

The second ionization energy is much greater than the first because the second electron being removed is in a lower energy level, closer to the nucleus. The electrons in lower energy levels are held more tightly by the protons in the nucleus, thus requiring more energy to remove.

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Most popular questions from this chapter

How much energy, in joules, must be absorbed to convert to \(\mathrm{Na}^{+}\) all the atoms present in \(1.00 \mathrm{mg}\) of gaseous Na? The first ionization energy of Na is \(495.8 \mathrm{kJ} / \mathrm{mol}.\)

Answer each of the following questions: (a) Which of the elements \(P, A s,\) and \(S\) has the largest atomic radius? (b) Which of the following has the smallest radius: \(\mathrm{Xe}, \mathrm{O}^{2-}, \mathrm{N}^{3-},\) or \(\mathrm{F}^{-} ?\) (c) Which should have the largest difference between the first and second ionization energy: \(\mathrm{Al}, \mathrm{Si}, \mathrm{P},\) or \(\mathrm{Cl} ?\) (d) Which has the largest ionization energy: \(\mathrm{C}, \mathrm{Si}\), or \(\mathrm{Sn}\) ? (e) Which has the largest electron affinity: \(\mathrm{Na}, \mathrm{B}\) \(\mathrm{Al},\) or \(\mathrm{C} ?\)

The most negative electron affinity of the following elements is that of (a) \(\mathrm{Br} ;\) (b) \(\mathrm{Sn} ;\) (c) \(\mathrm{Ba} ;\) (d) Li.

In your own words, define the following terms: (a) isoelectronic; (b) valence- shell electrons; (c) metal; (d) nonmetal; (e) metalloid.

Which of the following species would you expect tobe diamagnetic and which paramagnetic? (a) \(\mathrm{K}^{+}=\) (b) \(\mathrm{Cr}^{3+} ;\) (c) \(\mathrm{Zn}^{2+} ;\) (d) \(\mathrm{Cd} ;\) (e) \(\mathrm{Co}^{3+} ;\) (f) \(\mathrm{Sn}^{2+} ;\) (g) Br.
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