Question

In which of the following species, each atom carries same number of lone pair of electrons on it? (1) \(\mathrm{XeO}_{4}{ }^{2-}\) (2) \(\mathrm{Xe} \mathrm{F}_{2}\) (3) \(\mathrm{XeO}_{6}{ }^{4-}\) (4) \(\mathrm{O}_{3}\)

Short answer

\(\mathrm{XeF}_{2}\)

Step-by-step solution

- Understand the given species

Examine the molecular structure of each species to determine the number of lone pairs on each atom.

- Determine the electron count for each species

Calculate the total number of electrons and distribute them according to the molecular structure.

- Examine each species individually

Analyze each molecule to check if all the atoms have the same number of lone pairs. For example:- For \(\mathrm{XeO}_{4}{ }^{2-}\), Xe has 2 lone pairs, and O atoms have 2 lone pairs each.- For \(\mathrm{Xe} \mathrm{F}_{2}\), Xe has 3 lone pairs, and F atoms have 3 lone pairs each.- For \(\mathrm{XeO}_{6}{ }^{4-}\), Xe has no lone pairs, and O atoms have 3 lone pairs each.- For \(\mathrm{O}_{3}\), central O atom has 1 lone pair, and terminal O atoms have 2 lone pairs each.

- Identify the species with equal lone pairs on each atom

Identify the species where each atom in the molecule carries the same number of lone pairs.

- Conclusion

Based on the analysis, determine the correct answer.

Key concepts

molecular structure

Understanding molecular structure is fundamental to grasping how molecules interact and behave. A molecule's structure determines the arrangement of its atoms and the distribution of electrons.

Consider the molecular structure of \(\text{XeF}_{2}\):
- **Xenon** (Xe) acts as the central atom.
- **Fluorine** (F) atoms attach to the xenon.

The shape of \(\text{XeF}_{2}\) is linear due to pairs of non-bonding electrons (lone pairs) that repel each other, forcing the bonded atoms to align opposite each other.
Similarly, in \(\text{XeO}_{4}^{2-}\), the central xenon atom binds to four oxygen atoms, forming a tetrahedral structure.

Analyzing molecular structure helps in predicting:
- **Geometry**
- **Bond angles**
- **Electron distribution**
All of which are essential for understanding reactions and properties.

electron distribution

Electron distribution refers to how electrons are arranged in atoms and molecules. This is vital because electrons influence chemical properties and reactivity.

For example, let's look at one of the species in the exercise, \(\text{XeO}_{4}^{2-}\). Each **oxygen (O) atom** has 8 valence electrons:
- 6 from itself
- 2 more due to the negative charge

Now, \(\text{Xe}\) has 8 valence electrons:
- 3 are used to form bonds with the oxygen atoms
- Additionally, Xe has 2 lone pairs< br>
Distributing the electrons correctly enables us to draw accurate Lewis structures. This helps in:
- Predicting with which atoms molecules will form bonds
- Determine their shapes
- And chemical behavior

xenon compounds

Xenon compounds are fascinating as they showcase the versatility of a noble gas in forming stable molecules. Noble gases are typically unreactive due to having full valence shells. However, xenon can form compounds with various elements.

**Common xenon compounds** from the exercise include:
\(\text{XeF}_{2}\):
- Linear shape
- Xenon (Xe) has 3 lone pairs
- Each fluorine (F) atom has 3 lone pairs

\(\text{XeO}_{4}^{2-}\):
- Tetrahedral shape
- Xenon (Xe) has 2 lone pairs
- Each oxygen (O) atom has 2 lone pairs

In \(\text{XeO}_{6}^{4-}\):
- Octahedral shape
- Xenon (Xe) doesn't have any lone pairs
- Each oxygen (O) atom has 3 lone pairs

These structures highlight how xenon can utilize its d-orbitals to expand its valence shell beyond the octet rule, enabling the formation of such diverse and structurally rich compounds.