Question

Draw a plausible Lewis structure for the following series of molecules and ions: \((a) \operatorname{SiF}_{6}^{2-} ;\) (b) \(\mathrm{PF}_{5} ;\) (c) \(\mathrm{SF}_{4}\); (d) \(\mathrm{XeF}_{4}\). Describe the electron group geometry and molecular structure of these species.

Short answer

The Lewis structures are drawn for \(\operatorname{SiF}_{6}^{2-}\), \(PF_{5}\), \(SF_{4}\), and \(XeF_{4}\) according to the octet rule while considering the formal charge. The electron density around the central atoms gives the following geometries: octahedral for \(\operatorname{SiF}_{6}^{2-}\), trigonal bipyramidal for \(PF_{5}\), see-saw for \(SF_{4}\), and square planar for \(XeF_{4}\). As we progress down in the series, the electron group geometry remains the same - octahedral, but the molecular structure changes due to the presence of lone pairs of electrons.

Step-by-step solution

Draw Lewis Structure for \(\operatorname{SiF}_{6}^{2-}\)

Count total number of valence electrons. Silicon has 4 valence electrons and each Fluorine has 7. Add 2 for the 2- charge. So the total is 4+(6*7)+2=46. Draw Silicon in the center and six Fluorine atoms around it. Single bond each Fluorine atom to the Silicon and fill the rest electrons on Fluorine till it completes octet rule. Electron pairs are arranged around silicon in an octahedral.

Draw Lewis Structure for \(\mathrm{PF}_{5}\)

Count total number of valence electrons. Phosphorus has 5 and each Fluorine has 7. So the total is 5+(5*7)=40. Place the P in the center and five Fluorine atoms around it. Single bond each Fluorine atom to the Phosphorus. Fill rest of electrons on Fluorine till it completes octet rule. Electron pairs are arranged around phosphorous in a trigonal bipyramidal structure.

Draw Lewis Structure for \(\mathrm{SF}_{4}\)

Count total number of valence electrons. Sulfur has 6 and each Fluorine has 7. So the total is 6+(4*7)=34. Place the Sulfur in center and four Fluorine atoms around it. Single bond each Fluorine atom to the Sulfur. Electron pairs are arranged around sulfur in a see-saw structure.

Draw Lewis Structure for \( \mathrm{XeF}_{4}\)

Count total number of valence electrons. Xenon has 8 and each Fluorine has 7. So the total is 8+(4*7)=36. Place the Xenon in center and four Fluorine atoms around it. Single bond each Fluorine atom to the Xenon. Two remaining pairs of electrons on Xenon atom. Electron pairs are arranged around Xenon in a square planar structure.

Key concepts

Valence Electrons

Valence electrons play a crucial role in the formation of chemical bonds and the creation of molecules. These are the electrons located in the outermost shell of an atom, and they are primarily involved in forming bonds with other atoms. For example, the Silicon atom in \(\operatorname{SiF}_{6}^{2-}\),has 4 valence electrons.
Each surrounding Fluorine atom has 7 valence electrons, resulting in a total of 46 valence electrons once the additional 2 electrons from the negative charge are included.

• They determine how atoms interact chemically with each other.
• The Lewis structure is used to depict these interactions through dots and lines.
• Understanding the number of valence electrons helps in predicting bonding patterns and the stability of molecules.

Recognizing the number of valence electrons in atoms like Phosphorus, Sulfur, and Xenon helps us form accurate molecular structures like in \(\mathrm{PF}_5\), \(\mathrm{SF}_4\), and \(\mathrm{XeF}_4\).

Molecular Structure

Molecular structure refers to the three-dimensional arrangement of atoms in a molecule. It is how individual atoms are spatially oriented in a molecule. It's crucial to understand molecular structure to comprehend the shape and reactivity of the molecule.
For instance:

• In \(\mathrm{PF}_5\), Phosphorus occupies the central position, forming a trigonal bipyramidal structure with its five Fluorine atoms.
• In \(\mathrm{SF}_4\), Sulfur sits at the center with a see-saw geometry due to its four bonded Fluorine atoms.
• The geometry profoundly influences properties like polarity and intermolecular forces.

Understanding this concept can help predict how molecules will interact with one another. It guides chemists in determining molecular reactions and designing compounds with desired properties.

Electron Geometry

Electron geometry is the spatial arrangement of all pairs of electrons around the central atom, including both bond pairs and lone pairs. This geometry serves as the blueprint for predicting the shape of the molecule.By considering both bonded atoms and lone pairs:

• In\(\operatorname{SiF}_{6}^{2-}\),the electron geometry is octahedral as it encompasses six pairs of bonded electrons.
• For \(\mathrm{XeF}_4\), the electron geometry remains square planar due to four bonded atoms and two lone pairs on Xenon.

Electron geometry lays the groundwork for determining the molecular shape by influencing the angles and symmetry of the molecule.
Even with different molecular shapes, the electron geometry provides consistency in understanding molecular interactions.

Octet Rule

The octet rule is a fundamental chemical guideline stating that atoms tend to form bonds until they achieve a valence shell similar to that of noble gases, which typically means eight electrons.This rule explains the structure and bonding in molecules:

• Fluorine, for instance, adheres to the octet rule by forming bonds until achieving 8 electrons in its outer shell.
• In sulfur hexafluoride \(\mathrm{SF}_4\), the Sulfur atom deviates from the octet rule due to its larger size, allowing more than eight valence electrons.

This flexibility is crucial for understanding the limits and exceptions of the octet rule in real-world molecules.
Despite its limits, the octet rule remains a helpful tool in predicting bonding patterns and the stability of a given compound. In certain situations, like with Xenon and Phosphorus, extensions or exceptions to the rule are necessary for accommodating additional electrons.