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Chapter 8: Problem 80

Which of the following molecules or ions contain polar bonds: (a) \(\mathrm{P}_{4}\), (b) \(\mathrm{H}_{2} \mathrm{~S}\), (c) \(\mathrm{NO}_{2}^{-}\), (d) \(\mathrm{S}_{2}{ }^{2-}\) ?

Short Answer

Expert verified
Among the given molecules or ions, only \(\mathrm{NO}_{2}^{-}\) contains polar bonds. The bonds in \(\mathrm{P}_{4}\), \(\mathrm{H}_{2} \mathrm{~S}\), and \(\mathrm{S}_{2}{ }^{2-}\) are non-polar.

Step by step solution

01

(a) Analyzing \(\mathrm{P}_{4}\) bond polarity

In \(\mathrm{P}_{4}\), there are only phosphorus atoms bonded together. Since all these atoms have the same electronegativity, the bond is non-polar.
02

(b) Analyzing \(\mathrm{H}_{2} \mathrm{~S}\) bond polarity

For \(\mathrm{H}_{2} \mathrm{~S}\), we have hydrogen atoms bonded to a sulfur atom. The electronegativity values for hydrogen and sulfur are roughly 2.2 and 2.6, respectively. The difference between the electronegativity values is 0.4 which is less than 0.5. Hence, the \(\mathrm{H-S}\) bond is considered non-polar.
03

(c) Analyzing \(\mathrm{NO}_{2}^{-}\) bond polarity

In the case of \(\mathrm{NO}_{2}^{-}\), we have nitrogen bonded to two oxygen atoms. The electronegativity values for nitrogen and oxygen are 3.0 and 3.5, respectively. The difference between their electronegativity values is 0.5, which falls on the threshold for polarity. Consequently, the \(\mathrm{N-O}\) bond is considered polar.
04

(d) Analyzing \(\mathrm{S}_{2}{ }^{2-}\) bond polarity

For \(\mathrm{S}_{2} { }^{2-}\), we have two sulfur atoms bonded together. Since both atoms have the same electronegativity, the bond is non-polar. #Summary# In summary, only the \(\mathrm{NO}_{2}^{-}\) molecule contains polar bonds. The other molecules or ions have non-polar bonds.

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Most popular questions from this chapter

(a) Write one or more appropriate Lewis structures for the nitrite ion, \(\mathrm{NO}_{2}^{-} .\) (b) With what allotrope of oxygen is it isoelectronic? (c) What would you predict for the lengths of the bonds in \(\mathrm{NO}_{2}^{-}\) relative to \(\mathrm{N}-\mathrm{O}\) single bonds?

In the following pairs of binary compounds determine which one is a molecular substance and which one is an ionic substance. Use the appropriate naming convention (for ionic or molecular substances) to assign a name to each compound: (a) \(\mathrm{TiCl}_{4}\) and \(\mathrm{CaF}_{2}\), (b) \(\mathrm{ClF}_{3}\) and \(\mathrm{VF}_{3}\), (c) \(\mathrm{SbCl}_{5}\) and \(\mathrm{AlF}_{3}\).

(a) Construct a Lewis structure for hydrogen peroxide, \(\mathrm{H}_{2} \mathrm{O}_{2}\), in which each atom achieves an octet of electrons. (b) Do you expect the \(\mathrm{O}-\mathrm{O}\) bond in \(\mathrm{H}_{2} \mathrm{O}_{2}\) to be longer or shorter than the \(\mathrm{O}-\mathrm{O}\) bond in \(\mathrm{O}_{2}\) ?

The lattice energies of \(\mathrm{KBr}\) and \(\mathrm{CsCl}\) are nearly equal (Table 8.2). What can you conclude from this observation?

Acetylene \(\left(\mathrm{C}_{2} \mathrm{H}_{2}\right)\) and nitrogen \(\left(\mathrm{N}_{2}\right)\) both contain a triple bond, but they differ greatly in their chemical properties. (a) Write the Lewis structures for the two substances. (b) By referring to Appendix \(C\), look up the enthalpies of formation of acetylene and nitrogen and compare their reactivities. (c) Write balanced chemical equations for the complete oxidation of \(\mathrm{N}_{2}\) to form \(\mathrm{N}_{2} \mathrm{O}_{5}(g)\) and of acetylene to form \(\mathrm{CO}_{2}(g)\) and \(\mathrm{H}_{2} \mathrm{O}(g)\). (d) Calculate the enthalpy of oxidation per mole of \(\mathrm{N}_{2}\) and \(\mathrm{C}_{2} \mathrm{H}_{2}\) (the enthalpy of formation of \(\mathrm{N}_{2} \mathrm{O}_{5}(\mathrm{~g})\) is \(11.30 \mathrm{~kJ} / \mathrm{mol}\) ). How do these comparative values relate to your response to part (b)? Both \(\mathrm{N}_{2}\) and \(\mathrm{C}_{2} \mathrm{H}_{2}\) possess triple bonds with quite high bond enthalpies (Table 8.4). What aspect of chemical bonding in these molecules or
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