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Chapter 7: Problem 43

(a) What is the general relationship between the size of an atom and its first ionization energy? (b) Which element in the periodic table has the largest ionization energy? Which has the smallest?

Short Answer

Expert verified
(a) The general relationship between the size of an atom and its first ionization energy is an inverse relationship: as atomic size increases, ionization energy decreases. This is due to larger atoms having their outer electrons farther from the nucleus, which makes it easier to remove these electrons with less energy required. (b) In the periodic table, Helium (He) has the largest ionization energy, and Francium (Fr) has the smallest ionization energy.

Step by step solution

01

Understanding Ionization Energy

Ionization energy is the amount of energy required to remove an electron from a neutral atom in the gas phase. It is usually expressed in electron volts (eV) or kilojoules per mole (kJ/mol). Ionization energy generally increases from left to right across a period and decreases from top to bottom within a group in the periodic table. This is because as atoms get smaller, their nuclear charge increases, making it more difficult to remove an electron due to increased attraction.
02

Relationship Between Atomic Size and Ionization Energy

In general, there is an inverse relationship between atomic size and ionization energy. As the atomic size increases, the ionization energy decreases. This is because larger atoms have their outer electrons farther away from the nucleus, and thus, the nucleus has less of a "hold" on them, making it easier to remove these electrons with less energy required.
03

Finding the Element with the Largest Ionization Energy

The element with the largest ionization energy is the one in the top right corner of the periodic table, excluding the noble gases. This is because elements in the top right have smaller atomic sizes and greater nuclear charges, making it more difficult to remove an electron. In this case, the element with the largest ionization energy is Helium (He).
04

Finding the Element with the Smallest Ionization Energy

The element with the smallest ionization energy is the one in the bottom left corner of the periodic table. This is because elements in the bottom left have larger atomic sizes and smaller nuclear charges, making it easier to remove an electron. In this case, the element with the smallest ionization energy is Francium (Fr). Therefore, the general relationship between the size of an atom and its first ionization energy is that as atomic size increases, ionization energy decreases. The element with the largest ionization energy in the periodic table is Helium (He), and the one with the smallest ionization energy is Francium (Fr).

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Most popular questions from this chapter

Although the electron affinity of bromine is a negative quantity, it is positive for \(\mathrm{Kr}\). Use the electron configurations of the two elements to explain the difference.

The electron affinities, in \(\mathrm{kJ} / \mathrm{mol}\), for the group \(1 \mathrm{~B}\) and group \(2 \mathrm{~B}\) metals are $$ \begin{array}{|c|c|} \hline \mathrm{Cu} & \mathrm{Zn} \\ -119 & >0 \\\ \hline \mathrm{Ag} & \mathrm{Cd} \\ -126 & >0 \\ \hline \mathrm{Au} & \mathrm{Hg} \\ -223 & >0 \\ \hline \end{array} $$ (a) Why are the electron affinities of the group \(2 \mathrm{~B}\) elements greater than zero? (b) Why do the electron affinities of the group \(1 \mathrm{~B}\) elements become more negative as we move down the group? [Hint: Examine the trends in the electron affinity of other groups as we proceed down the periodic table. \(]\)

Find three examples of ions in the periodic table that have an electron configuration of \(n d^{8}(n=3,4,5 \ldots) .\)

(a) Write the electron configuration for \(\mathrm{Li}\), and estimate the effective nuclear charge experienced by the valence electron. (b) The energy of an electron in a one-electron atom or ion equals \(\left(-2.18 \times 10^{-18} \mathrm{~J}\right)\left(\frac{Z^{2}}{n^{2}}\right)\) where \(Z\) is the nuclear charge and \(n\) is the principal quantum number of the electron. Estimate the first ionization energy of Li. (c) Compare the result of your calculation with the value reported in Table 7.4 and explain the difference. (d) What value of the effective nuclear charge gives the proper value for the ionization energy? Does this agree with your explanation in \((\mathrm{c}) ?\)

(a) What is an isoelectronic series? (b) Which neutral atom is isoelectronic with each of the following ions: \(\mathrm{Ga}^{3+}, \mathrm{Zr}^{4+}\) \(\mathrm{Mn}^{7+}, \Gamma, \mathrm{Pb}^{2+} ?\)
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