Question

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

Short answer

The Lewis structures for the given molecules or ions are as follows: (a) \(\mathrm{NH}_{4}^{+}\) obeys the octet rule: ``` H | H - N - H | H ``` (b) \(\mathrm{SCN}^{-}\) obeys the octet rule: ``` S = C - N ``` (c) \(\mathrm{PCl}_{3}\) obeys the octet rule: ``` Cl | Cl - P - Cl ``` (d) \(\mathrm{TeF}_{4}\) does not obey the octet rule: ``` F | F - Te - F | F ``` (e) \(\mathrm{XeF}_{2}\) does not obey the octet rule: ``` F | Xe - F ``` In summary, \(\mathrm{NH}_{4}^{+}\), \(\mathrm{SCN}^{-}\), and \(\mathrm{PCl}_{3}\) obey the octet rule, while \(\mathrm{TeF}_{4}\) and \(\mathrm{XeF}_{2}\) do not.

Step-by-step solution

Total number of valence electrons

In \(\mathrm{NH}_{4}^{+}\), nitrogen has 5 valence electrons, hydrogen has 1 valence electron, and the positive charge indicates that there is one less electron. So, the total number of valence electrons is: \((5 * 1) + (4 * 1) - 1 = 8\).

Draw the Lewis structure

We start by placing nitrogen in the center with single bonds to each hydrogen atom. To complete the octet of nitrogen, we use the remaining electrons to form lone pairs. Remaining electrons are 0, which means that the chosen structure fulfills the octet rule. ``` H | H - N - H | H ```

Does it obey the octet rule?

Yes, \(\mathrm{NH}_{4}^{+}\) obeys the octet rule. ##(b) \(\mathrm{SCN}^{-}\)##

Total number of valence electrons

In \(\mathrm{SCN}^{-}\), sulfur has 6 valence electrons, carbon has 4 valence electrons, nitrogen has 5 valence electrons, and the negative charge indicates that there is one additional electron: \((6 * 1) + (4 * 1) + (5 * 1) + 1 = 16\).

Draw the Lewis structure

We place carbon in the center and form single bonds with both sulfur and nitrogen. We then complete the octet with lone pairs for sulfur and nitrogen. ``` S - C - N ``` To satisfy the octet rule for carbon, we need to add multiple bonds, one double bond can be added either to sulfur or nitrogen. ``` S = C - N ```

Does it obey the octet rule?

Yes, \(\mathrm{SCN}^{-}\) obeys the octet rule. ##(c) \(\mathrm{PCl}_{3}\)##

Total number of valence electrons

In \(\mathrm{PCl}_{3}\), phosphorus has 5 valence electrons, and chlorine has 7 valence electrons. The total number of valence electrons is: \((5 * 1) + (7 * 3) = 26\).

Draw the Lewis structure

We place phosphorus in the center and form single bonds with each of the three chlorine atoms. To complete the octet for each chlorine, we use the remaining electrons to form lone pairs. ``` Cl | Cl - P - Cl ```

Does it obey the octet rule?

Yes, \(\mathrm{PCl}_{3}\) obeys the octet rule. ##(d) \(\mathrm{TeF}_{4}\)##

Total number of valence electrons

In \(\mathrm{TeF}_{4}\), tellurium has 6 valence electrons, and fluorine has 7 valence electrons. The total number of valence electrons is: \((6 * 1) + (7 * 4) = 34\).

Draw the Lewis structure

We place tellurium in the center and form single bonds with each of the four fluorine atoms. To complete the octet for each fluorine, we use the remaining electrons to form lone pairs. ``` F | F - Te - F | F ``` With remaining two electrons, we form a lone pair on the tellurium.

Does it obey the octet rule?

No, \(\mathrm{TeF}_{4}\) does not obey the octet rule, since tellurium has 10 electrons around it. ##(e) \(\mathrm{XeF}_{2}\)##

Total number of valence electrons

In \(\mathrm{XeF}_{2}\), xenon has 8 valence electrons, and fluorine has 7 valence electrons. The total number of valence electrons is: \((8 * 1) + (7 * 2) = 22\).

Draw the Lewis structure

We place xenon in the center and form single bonds with each of the two fluorine atoms. To complete the octet for each fluorine, we use the remaining electrons to form lone pairs. ``` F | Xe - F ``` With remaining four electrons, we form two lone pairs on the xenon.

Does it obey the octet rule?

No, \(\mathrm{XeF}_{2}\) does not obey the octet rule, since xenon has 12 electrons around it. In summary, \(\mathrm{NH}_{4}^{+}\), \(\mathrm{SCN}^{-}\), and \(\mathrm{PCl}_{3}\) obey the octet rule, while \(\mathrm{TeF}_{4}\) and \(\mathrm{XeF}_{2}\) do not.

Key concepts

Octet Rule

The octet rule is a fundamental principle in chemistry that explains how atoms bond to achieve more stable arrangements. It states that atoms tend to form compounds in ways that allow them to have eight electrons in their valence shell, emulating the electron configuration of noble gases.

This rule is key to understanding why certain atoms interact and bond the way they do. For example, in the molecule ammonium (\(NH_{4}^{+}\)), nitrogen achieves an octet by sharing its electrons with four hydrogen atoms. On the other hand, tellurium in \(TeF_{4}\) and xenon in \(XeF_{2}\) exceed the octet, which is possible because these elements have access to d-orbitals that can accommodate more than eight electrons. Such exceptions typically occur with elements located in the third period or below in the periodic table.

Valence Electrons

Valence electrons are the electrons found in the outermost shell of an atom that determine its chemical properties and bonding behavior. They are essentially the 'currency' for atomic interactions, as they are the electrons involved in forming bonds. In the solutions provided, the total number of valence electrons in a molecule dictates how the atoms can be bonded together to form a stable structure.

For instance, phosphorus in \(PCl_3\) has five valence electrons and needs three more to complete its octet. It achieves this by forming single bonds with three chlorine atoms, each providing one electron. Valence electrons are depicted in Lewis structures as dots around the atomic symbols, showcasing the electron-sharing or lone pairs.

Electron Dot Structure

Electron dot structures, commonly known as Lewis structures, are a visual representation of the bonding between atoms and the lone pair electrons. The dots signify valence electrons, while lines represent shared pairs of electrons, or bonds. The goal is to display the arrangement of electrons that results in a stable molecular structure.

While creating Lewis structures, each atom's need to fulfill the octet rule (or duet rule for hydrogen) dictates the distribution of electrons. In the case of simple molecules such as \(NH_{4}^{+}\), drawing the Lewis structure is straightforward, but for compounds like \(SCN^{-}\), deciding where to place a double bond to satisfy carbon's octet can be less clear. Ensuring that all valence electrons are accounted for and that each atom (with some exceptions) achieves an octet is essential in constructing accurate Lewis structures.