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Chapter 9: Problem 55

Why does the octet rule not hold for many compounds containing elements in the third period of the periodic table and beyond?

Short Answer

Expert verified
The octet rule does not hold for many compounds containing elements in the third period of the periodic table and beyond due to the presence of d orbitals in their third energy level. These d orbitals allow the accommodation of more than eight valence electrons, violating the octet rule.

Step by step solution

01

Understanding the Octet Rule

The octet rule states that atoms tend to combine in such a way that they each have eight electrons in their valence shells, giving them the same electronic configuration as a noble gas. This rule is based on the principle that having a full valence shell is the most stable electron configuration for an atom.
02

Identifying why the Octet Rule doesn't always hold

Periods 1 and 2 in the periodic table only have s and p orbitals available in their outermost energy level, limiting them to a maximum of eight valence electrons. However, starting from period 3, elements also have d orbitals available in their third energy level. These d orbitals can accomodate more electrons, allowing these elements to have more than eight electrons in their valence shells. This explains why the octet rule is violated by elements in the third period and beyond.
03

Specific example: Phosphorus

Take phosphorus for example, an element from the third period of the periodic table. Phosphorus has five valence electrons in its outer shell. However, due to the presence of empty 3d orbitals, it can expand its valence shell to accommodate more than eight electrons. A common compound that exemplifies this is phosphorus pentachloride (PCl5), where phosphorus is surrounded by five chlorine atoms, violating the octet rule.

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Most popular questions from this chapter

Draw reasonable resonance structures for these (a) \(\mathrm{HSO}_{4}^{-},(\mathrm{b}) \mathrm{SO}_{4}^{2-},(\mathrm{c}) \mathrm{HSO}_{3}^{-}\), sulfur-containing ions: (d) \(\mathrm{SO}_{3}^{2-}\).

Which of these molecules or ions are isoelectronic: \(\mathrm{NH}_{4}^{+}, \mathrm{C}_{6} \mathrm{H}_{6}, \mathrm{CO}, \mathrm{CH}_{4}, \mathrm{~N}_{2}, \mathrm{~B}_{3} \mathrm{~N}_{3} \mathrm{H}_{6} ?\)

Draw three reasonable resonance structures for the \(\mathrm{OCN}^{-}\) ion. Show formal charges.

Using this information: $$ \begin{array}{rr} \mathrm{C}(s) \longrightarrow \mathrm{C}(g) & \Delta H_{\mathrm{rxn}}^{\circ}=716 \mathrm{~kJ} / \mathrm{mol} \\ 2 \mathrm{H}_{2}(g) \longrightarrow 4 \mathrm{H}(g) & \Delta H_{\mathrm{rxn}}^{\circ}=872.8 \mathrm{~kJ} / \mathrm{mol} \end{array} $$ and the fact that the average \(\mathrm{C}-\mathrm{H}\) bond enthalpy is \(414 \mathrm{~kJ} / \mathrm{mol}\), estimate the standard enthalpy of formation of methane \(\left(\mathrm{CH}_{4}\right)\).

Draw three resonance structures for the chlorate ion, \(\mathrm{ClO}_{3}^{-}\). Show formal charges.
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