Question

Which of the following statements is(are) true? Correct the false statements. a. It is impossible to satisfy the octet rule for all atoms in \(\mathrm{XeF}_{2}\) b. Because \(\mathrm{SF}_{4}\) exists, \(\mathrm{OF}_{4}\) should also exist because oxygen is in the same family as sulfur. c. The bond in \(\mathrm{NO}^{+}\) should be stronger than the bond in \(\mathrm{NO}^{-}\) d. As predicted from the two Lewis structures for ozone, one oxygen-oxygen bond is stronger than the other oxygenoxygen bond.

Short answer

The corrected statements are: a. The octet rule is satisfied for all atoms in \(\mathrm{XeF}_{2}\). b. The existence of \(\mathrm{SF}_{4}\) does not guarantee the existence of \(\mathrm{OF}_{4}\) as oxygen has a higher electronegativity and typically forms a maximum of two bonds. c. The bond in \(\mathrm{NO}^{-}\) should be stronger than the bond in \(\mathrm{NO}^{+}\) as it contains a triple bond, while \(\mathrm{NO}^{+}\) contains a double bond. d. In ozone, both oxygen-oxygen bonds are equivalent in strength due to resonance, causing them to be a hybrid of single and double bonds.

Step-by-step solution

a. Analyzing the octet rule in \(\mathrm{XeF}_{2}\)

First, we need to draw the Lewis structure of \(\mathrm{XeF}_{2}\). Xenon (Xe) has 8 valence electrons, and Fluorine (F) has 7. Since there are two Fluorine atoms, the total number of valence electrons in the molecule is 8 + 2 * 7 = 22. In the Lewis structure, there will be two single bonds between Xe and two F atoms, which will satisfy the octet rule for the two F atoms. The remaining 18 electrons (3 lone pairs) will be placed around the Xe atom, satisfying the octet rule for the Xe as well. Thus, this statement is false since the octet rule is satisfied for all atoms in \(\mathrm{XeF}_{2}\).

b. Comparing \(\mathrm{SF}_{4}\) and \(\mathrm{OF}_{4}\)

Sulfur and oxygen belong to the same family (Group 16); however, this does not guarantee that \(\mathrm{OF}_{4}\) will exist if \(\mathrm{SF}_{4}\) does. Oxygen is more electronegative than sulfur, making it less likely to form a molecule like \(\mathrm{OF}_{4}\). Oxygen typically forms a maximum of two bonds due to its high electronegativity. Therefore, the statement is false.

c. Comparing bond strengths in \(\mathrm{NO}^{+}\) and \(\mathrm{NO}^{-}\)

In \(\mathrm{NO}^{+}\), the nitrogen and oxygen atoms share two electrons, forming a double bond, while in \(\mathrm{NO}^{-}\), they share three electrons, resulting in a triple bond. Triple bonds are generally stronger than double bonds because they involve more electron sharing between the atoms. Thus, the statement is false since the bond in \(\mathrm{NO}^{-}\) should be stronger than the bond in \(\mathrm{NO}^{+}\).

d. Analyzing the oxygen-oxygen bonds in ozone

Ozone, \(\mathrm{O}_{3}\), has two resonance structures with one double bond and one single bond between the oxygen atoms. However, due to resonance, the actual structure of ozone is a hybrid of these two resonance structures. This means that both oxygen-oxygen bonds are equivalent in terms of bond order and strength. Consequently, the statement is false because the two oxygen-oxygen bonds in ozone are not different in strength since they are both a combination of single and double bonds due to resonance. In summary, all four statements given are false.

Key concepts

Octet Rule

The octet rule is a fundamental concept in chemistry that suggests atoms tend to bond in such a way that each atom has eight electrons in its valence shell, leading to more stable electron configurations. This can be achieved through the sharing, losing, or gaining of electrons.
For example, in the molecule \( \text{XeF}_2 \), xenon (Xe) and fluorine (F) atoms bond so that the octet rule is satisfied. This involves sharing electrons between Xe and the two F atoms. While xenon is an exception to the usual octet rule due to its larger size and capacity to accommodate more electrons, in \( \text{XeF}_2 \), it holds 18 valence electrons, thus having a "super octet."
Understanding the octet rule can help predict how molecules will form and what their structures might look like when drawing Lewis structures.

Lewis Structures

Lewis structures are a simple yet effective way to represent the bonding and electron distribution in molecules. They involve drawing dots around atoms to illustrate valence electrons and using lines to show bonds between atoms.
To draw a Lewis structure, follow these steps:

• Count the total number of valence electrons available in the molecule.
• Connect the atoms with single bonds first.
• Distribute remaining electrons to satisfy the octet rule for each atom, starting with the most electronegative elements.

In the case of \( \text{XeF}_2 \), the Lewis structure will reveal that xenon shares two pairs of electrons with fluorine atoms, with three lone pairs remaining on xenon. This clear visualization aids in understanding molecular geometry and reactivity.

Resonance

Resonance occurs when two or more valid Lewis structures can be drawn for a molecule or ion. These different structures, called resonance forms or structures, differ only in the distribution of electrons, not the arrangement of atoms. They contribute to a resonance hybrid, which is the true structure that is an average of all resonance forms.
Take ozone (\( \text{O}_3 \)), for instance. It has two resonance forms with varying locations of double and single oxygen-oxygen bonds. However, in reality, both bonds are equivalent in terms of strength as they are a blend of the resonance structures. This concept helps explain phenomena such as bond length and energy that single Lewis structures cannot accurately depict.

Electronegativity

Electronegativity refers to the ability of an atom to attract bonding electrons. It's a crucial factor in determining the type of bond that will form between atoms, whether it's ionic, covalent, or polar covalent.
For example, in comparing sulfur (S) and oxygen (O), oxygen is more electronegative. This means oxygen holds onto electrons more tightly than sulfur does. This increased electronegativity is the reason \( \text{OF}_4 \) is unlikely to exist as a stable molecule; oxygen's tendency to attract electrons limits its ability to form four bonds like sulfur in \( \text{SF}_4 \). Understanding electronegativity helps in predicting molecular polarities and reactivities.