Question

Draw the Lewis structures for the molecules \(\$ F_{4}\) and \(S F_{6} .\) Analyze each structure to determine whether the molecule is polar or nonpolar.

Short answer

The Lewis structure for \(F_4\) has a symmetrical tetrahedral shape with equal bond angles of 109.5˚. As all the surrounding atoms are the same (F), the molecule is nonpolar. Similarly, the Lewis structure for \(SF_6\) has a symmetrical octahedral shape with equal bond angles of 90˚ and 180˚. As all the surrounding atoms are also the same (F), this molecule is nonpolar as well. In conclusion, both \(F_4\) and \(SF_6\) molecules are nonpolar.

Step-by-step solution

Finding the total number of valence electrons for the molecules

In this step, we will find out the total number of valence electrons in the molecules \(F_4\) and \(SF_6\). Fluorine (F) belongs to Group 17, so it has 7 valence electrons. Sulfur (S) belongs to Group 16 and has 6 valence electrons. For \(F_4\): There are 4 Fluorine atoms in this molecule, so the total number of valence electrons = 4*7 = 28 For \(SF_6\): There is 1 Sulfur and 6 Fluorine atoms in this molecule, so the total number of valence electrons = 1*6 + 6*7 = 6 + 42 = 48

Drawing the Lewis structures

Now that we know the total number of valence electrons, we can draw the Lewis structures for both molecules. For \(F_4\): 1. Connect the 4 Fluorine atoms to each other in a square structure. 2. Distribute the remaining electrons to complete the octet of each atom (except Hydrogen). Each F atom already has 2 shared electrons and needs 6 more. For \(SF_6\): 1. Put Sulfur as the central atom and connect the 6 Fluorine atoms around it. 2. Distribute the remaining electrons to complete the octet of each atom (except Hydrogen). Sulfur is already sharing 12 electrons, and each F atom has 2 shared electrons and needs 6 more.

Determining the electron pair geometry and molecular geometry

Now, we need to determine the electron pair geometry and molecular geometry of the molecules to be able to understand if the molecules are polar or nonpolar. For \(F_4\): The electron pair geometry has 4 electron pair regions around the central atom. Therefore, the electron pair geometry is tetrahedral. Since all the bond angles are 109.5˚, the molecular geometry is also tetrahedral. For \(SF_6\): The electron pair geometry has 6 electron pair regions around the central atom. Therefore, the electron pair geometry is octahedral. Since all the bond angles are 90˚ and 180˚, the molecular geometry is also octahedral.

Determining if the molecule is polar or nonpolar based on the molecular geometry

By analyzing the molecular geometry, we can determine if the molecule is polar or nonpolar. For \(F_4\): The molecule has a symmetrical tetrahedral shape, and all the atoms surrounding the central atom (F) are the same. Therefore, the molecule is nonpolar. For \(SF_6\): The molecule has a symmetrical octahedral shape, and all the atoms surrounding the central atom (S) are the same. Therefore, the molecule is also nonpolar. In conclusion, both \(F_4\) and \(SF_6\) molecules are nonpolar.

Key concepts

Understanding Valence Electrons

Valence electrons are the outermost electrons of an atom and play a crucial role in chemical bonding.
These electrons determine how an atom will interact or bond with another atom.
For instance, fluorine (F) is in Group 17 of the periodic table and has 7 valence electrons, whereas sulfur (S) from Group 16 has 6 valence electrons.

• To calculate the total number of valence electrons in a molecule, multiply the number of each type of atom by its valence electrons and sum them up.
  • For \(F_4\), there are 4 fluorine atoms: \(4 \times 7 = 28\) valence electrons.
  • For \(SF_6\), there are 1 sulfur and 6 fluorine atoms: \(1 \times 6 + 6 \times 7 = 48\) valence electrons.

Valence electrons are fundamental in drawing Lewis structures and predicting molecular behavior.

Decoding Molecular Geometry

Molecular geometry refers to the three-dimensional arrangement of atoms in a molecule.
This geometry affects many physical and chemical properties, including potential polarity.
When predicting molecular geometry, it is essential to use the VSEPR (Valence Shell Electron Pair Repulsion) theory, which helps us understand the spatial orientation of electron pairs.

• For \(F_4\): The molecule forms a tetrahedral shape, a structure with a symmetrical distribution of electron pairs leading to identical bond angles of 109.5°.
• For \(SF_6\): The molecule forms an octahedral shape, characterized by bond angles of 90° and 180°.

This knowledge helps determine the molecule's potential behavior during interactions.

Polarity: Understanding Molecular Interactions

Polarity is a measure of how evenly electric charge is distributed across a molecule.
It significantly impacts a molecule’s interaction with other molecules, affecting properties like solubility and boiling point.
Molecular shape and symmetry often dictate polarity.

• In a molecule with a symmetrical shape, if atoms surrounding the central atom are similar, dipoles cancel out, rendering the molecule nonpolar.
• For \(F_4\): With all identical atoms in a tetrahedral configuration, the dipoles balance out, making it nonpolar.
• For \(SF_6\): The symmetrical octahedral shape ensures that the polar bonds cancel, resulting in a nonpolar molecule.

Nonpolarity in symmetrical molecules indicates uniform charge distribution.

Exploring Electron Pair Geometry

Electron pair geometry considers the spatial arrangement of all electron pairs (bonding and lone pairs) around the central atom.
It is a precursor for deducing molecular geometry as it provides insight into how electron repulsions shape the molecule.

• For \(F_4\): Electron pair geometry is described as tetrahedral since there are four bonding pairs and no lone pairs around the central atom.
• For \(SF_6\): It is octahedral due to the presence of six bonding pairs around the sulfur atom.

By studying electron pair geometry, we grasp how molecules form specific shapes essential in predicting chemical behavior.

The Significance of the Octet Rule

The octet rule is a guiding principle in chemistry that states atoms tend to form bonds until they are surrounded by 8 valence electrons.
Achieving an octet signifies stability for the atom, as is seen in noble gases.
While this rule guides the formation of Lewis structures, it also explains why certain molecules such as \(SF_6\) can "expand" their octet.

• For \(F_4\): Each fluorine atom completes its octet by sharing electrons with adjacent atoms.
• For \(SF_6\): Sulfur exceeds the octet by transferring electrons into available \(d\) orbitals, allowing it to bond with six fluorine atoms.

Although the octet rule is a significant guideline, it is not absolute, as seen in hypervalent molecules like \(SF_6\).