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Chapter 8: Problem 62

Draw the Lewis structures for each of the following molecules or ions. Which do not obey the octet rule? (a) \(\mathrm{NH}_{4}^{+}\), (b) \(\mathrm{SCN}^{-}\), (c) \(\mathrm{PCl}_{3}\), (d) \(\mathrm{TeF}_{4}\) (e) \(\mathrm{XeF}_{2}\).

Short Answer

Expert verified
The Lewis structures for the given molecules/ions are: (a) \(\mathrm{NH}_{4}^{+}\): H-N-H with an additional H atom connected to N above and below. (b) \(\mathrm{SCN}^{-}\): S-C鈮 (c) \(\mathrm{PCl}_{3}\): Cl-P-Cl with an additional Cl atom connected to P above. (d) \(\mathrm{TeF}_{4}\): F-Te-F with an additional F atom connected to Te above and below. (e) \(\mathrm{XeF}_{2}\): F-Xe, with an additional F atom connected to Xe below. All of them obey the octet rule; however, TeF4 and XeF2 expand their octet, meaning their central atoms can have more than 8 electrons.

Step by step solution

01

Determine the number of valence electrons

First, we need to find the number of valence electrons for each atom in the molecules/ions. The number of valence electrons is equal to the group number of the atom in the periodic table. In case of ions, add or subtract the charge to the total number of valence electrons. The number of valence electrons for each molecule or ion is as follows: (a) \(\mathrm{NH}_{4}^{+}\): N has 5 valence electrons and each of the 4 H atoms has 1. The positive charge means we subtract 1 electron, for a total of 8 valence electrons. (b) \(\mathrm{SCN}^{-}\): S has 6 valence electrons, C has 4, and N has 5. The negative charge means we add 1 electron, for a total of 16 valence electrons. (c) \(\mathrm{PCl}_{3}\): P has 5 valence electrons and each of the 3 Cl atoms has 7, for a total of 26 valence electrons. (d) \(\mathrm{TeF}_{4}\): Te has 6 valence electrons and each of the 4 F atoms has 7, for a total of 34 valence electrons. (e) \(\mathrm{XeF}_{2}\): Xe has 8 valence electrons and each of the 2 F atoms has 7, for a total of 22 valence electrons.
02

Draw the Lewis structures

To draw the Lewis structures, we need to connect the atoms by placing pairs of electrons between them and arranging the remaining electrons around the atoms to satisfy the octet rule whenever possible. (a) \(\mathrm{NH}_{4}^{+}\): Place N at the center, connecting it to four H atoms. The central N has a full octet. H | H-N-H | H (b) \(\mathrm{SCN}^{-}\): Place S at one end, N at the other end, and C in the middle. Connect S to C and C to N. Add the remaining valence electrons to satisfy the octet rule. S-C鈮 (c) \(\mathrm{PCl}_{3}\): Place P at the center and connect it to three Cl atoms. Satisfy the octet rule by adding the remaining valence electrons. Cl | Cl-P-Cl | (d) \(\mathrm{TeF}_{4}\): Place Te at the center and connect it to four F atoms. Because Te can expand its octet, we can satisfy the octet rule for all atoms. F F | | F-Te-F | (e) \(\mathrm{XeF}_{2}\): Place Xe at the center and connect it to two F atoms. Since Xe can expand its octet, we can add the remaining electrons to satisfy the octet rule for all atoms. F | F-Xe | F
03

Identify molecules/ions that do not obey the octet rule

From the Lewis structures above, we can see that all of the given molecules and ions obey the octet rule. However, some of them (TeF4 and XeF2) expand their octet, meaning they can have more than 8 electrons around the central atom.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Valence Electrons
Valence electrons are the outermost electrons of an atom and play a crucial role in chemical bonding. They determine how an atom can bond with others and are responsible for the chemical behavior of the element. To find the number of valence electrons, look at the group number in the Periodic Table. For example, carbon (C) is in Group 14, making it have 4 valence electrons. In ions, adjust the count of electrons based on the charge:
  • If the ion is negatively charged, add those electrons to the total count.
  • If it is positively charged, subtract them.
With valence electrons identified, you can start drawing Lewis Structures, representing these electrons as dots around the symbols of the elements.
Octet Rule
The Octet Rule is a fundamental concept in chemistry, stipulating that atoms tend to bond in such a way that they each have eight electrons in their valence shell, achieving a stable, noble gas configuration. This rule drives the bonding interactions in molecules:
  • Atoms will share, gain, or lose electrons to complete their octet.
  • Hydrogen is an exception, as it is stable with two electrons in its outer shell, resembling helium.
Drawing Lewis Structures helps visualize whether the octet rule is satisfied. In some cases, elements can "expand" their octet, especially those in period 3 or beyond on the periodic table. Such elements have d-orbitals available, allowing more than 8 electrons around them. Examples include TeF鈧� and XeF鈧�, which accept more electrons around the central atom, exceeding the octet for stability.
Molecular Geometry
Molecular Geometry refers to the three-dimensional arrangement of atoms within a molecule. It determines not only the shape but also the reactivity, polarity, phase of matter, and biological activity of the molecules. Visualizing geometry is crucial:
  • Linear, bent, trigonal pyramidal, and tetrahedral are some common shapes.
  • Lone pairs of electrons on the central atom can influence the geometry, often resulting in more complex shapes.
For example, in PCl鈧�, the shape is trigonal pyramidal due to lone pairs on the phosphorus, while XeF鈧� has a linear shape with five electron pairs around the xenon atom. Understanding these shapes can predict interaction with other molecules and the resultant properties.
Chemical Bonding
Chemical Bonding describes the forces holding atoms together within molecules, dictating how molecules interact with other substances. Two primary types of bonding are covalent and ionic. Understanding bonding is essential:
  • Covalent bonding involves sharing electron pairs between atoms, forming strong bonds within a molecule. This is typical in non-metals.
  • Ionic bonding occurs when electrons are transferred from one atom to another, creating positively and negatively charged ions held together by electrostatic forces.
Lewis Structures are a visual tool in depicting covalent bonds, showing how atoms share electrons to fulfill the octet rule. For example, in SCN鈦�, covalent bonds involve sharing electrons between S, C, and N, achieving stable configurations. Understanding bonding helps foresee molecule stability, reactivity, and the type of chemical reactions they may undergo.

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Most popular questions from this chapter

What is the most common exception to the octet rule? Give two examples.

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

(a) How does a polar molecule differ from a nonpolar one? (b) Atoms X and Y have different electronegativities. Will the diatomic molecule \(X-Y\) necessarily be polar? Explain. (c) What factors affect the size of the dipole moment of a diatomic molecule?

Average bond enthalpies are generally defined for gasphase molecules. Many substances are liquids in their standard state. thermochemical data from Appendix \(C\), calculate average bond enthalpies in the liquid state for the following bonds, and compare these values to the gas-phase values given in Table 8.4: (a) \(\mathrm{Br}-\mathrm{Br}\), from \(\mathrm{Br}_{2}(l) ;\) (b) \(\mathrm{C}-\mathrm{Cl}\), from \(\mathrm{CCl}_{4}(l) ;\) (c) \(\mathrm{O}-\mathrm{O}\), from \(\mathrm{H}_{2} \mathrm{O}_{2}(l)\) (assume that the \(\mathrm{O}-\mathrm{H}\) bond enthalpy is the same as in the gas phase). (d) What can you conclude about the process of breaking bonds in the liquid as compared to the gas phase? Explain the difference in the \(\Delta H\) values between the two phases.

The dipole moment and bond distance measured for the highly reactive gas phase OH molecule are \(1.78 \mathrm{D}\) and \(0.98 \AA\), respectively. (a) Given these values calculate the effective charges on the \(\mathrm{H}\) and \(\mathrm{O}\) atoms of the OH molecule in units of the electronic charges \(e\). (b) Is this bond more or less polar than the \(\mathrm{H}-\mathrm{Cl}\) bond in an \(\mathrm{HCl}\) molecule? (c) Is that what you would have expected based on electronegativities?
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