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Chapter 8: Problem 51

Use the third period of the periodic table as an example to illustrate the change in first ionization energies of the elements as we move from left to right. Explain the trend.

Short Answer

Expert verified
As we move from left to right across the third period of the periodic table, the first ionization energy increases due to increasing nuclear charge and decreasing atomic radius. However, exceptions to this trend occur due to the electron configuration of some elements (Al, S).

Step by step solution

01

Identify the Elements of the Third Period

The third period of the periodic table consists of the elements from sodium (Na) to argon (Ar). These elements are: Na, Mg, Al, Si, P, S, Cl, and Ar.
02

Define First Ionization Energy

The first ionization energy is the energy required to remove the most loosely bound electron, the valence electron, from a neutral atom in its gaseous phase.
03

Analyze the Trend

As we move from left to right across the third period of the periodic table, the first ionization energy generally increases. This trend is mainly due to two effects: nuclear charge and atomic radius. As we go across the period, the atomic number increases, which means an increase in nuclear charge. This greater nuclear charge pulls the electrons in the atom closer, reducing the atomic radius and thus making electrons harder to remove, thereby increasing ionization energy.
04

Consider Exceptions

Although the general trend is for the first ionization energy to increase as you move from left to right, there are some exceptions. For sodium (Na) to magnesium (Mg), and phosphorus (P) to sulfur (S), the first ionization energy increases. However, aluminum (Al) and silicon (Si) do not follow this pattern, as the first ionization energy of aluminum is less than that of magnesium. This also occurs with phosphorus and sulfur.
05

Explain the Exceptions

These exceptions occur because of the structure of the electron shells. For Al and S, an electron is being removed from a higher-energy subshell (3p), which is easier to remove than from a lower-energy subshell (3s). This explains why the first ionization energy decreases from Mg to Al and from P to S.
06

Conclusion

On the whole, as you move from left to right across the third period of the periodic table, the first ionization energy tends to increase due to increasing nuclear charge and decreasing atomic radius. However, there are exceptions due to the electron configuration of some elements.

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Most popular questions from this chapter

For each pair of elements listed here, give three properties that show their chemical similarity: (a) sodium and potassium and (b) chlorine and bromine.

Use the alkali metals and alkaline earth metals as examples to show how we can predict the chemical properties of elements simply from their electron configurations.

The first and second ionization energies of \(\mathrm{K}\) are \(419 \mathrm{~kJ} / \mathrm{mol}\) and \(3052 \mathrm{~kJ} / \mathrm{mol}\), and those of \(\mathrm{Ca}\) are \(590 \mathrm{~kJ} / \mathrm{mol}\) and \(1145 \mathrm{~kJ} / \mathrm{mol}\), respectively. Compare their values and comment on the differences.

You are given a dark shiny solid and asked to determine whether it is iodine or a metallic element. Suggest a nondestructive test that would enable you to arrive at the correct answer.

Which elements are more likely to form acidic oxides? basic oxides? amphoteric oxides?
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