Question

For each compound, draw the Lewis structure, determine the geometry using \(\mathrm{VSEPR}\) theory, determine whether the molecule is polar, identify the hybridization of all interior atoms, and make a sketch of the molecule, according to valence bond theory, showing orbital overlap. a. \(\mathrm{COF}_{2}(\) carbon is the central atom \()\) b. \(\mathrm{S}_{2} \mathrm{Cl}_{2}(\mathrm{ClSSCl})\) c. \(\mathrm{SF}_{4}\)

Short answer

COF2: Polar, trigonal planar, sp2 hybridized. S2Cl2: Polar, bent shape for each SCl2 unit, sp3 hybridized. SF4: Polar, see-saw shape, sp3d hybridization.

Step-by-step solution

Draw Lewis Structure for COF2

The first step is to draw the Lewis Structure for COF2. Carbon has 4 valence electrons, oxygen has 6, and each fluorine has 7. Carbon is the central atom. Single bonds are formed between carbon and oxygen and carbon and each fluorine. Ten electrons remain to fill the octets of oxygen and fluorine, resulting in a double bond with oxygen and completing the octets of the fluorine atoms. The carbon will have no lone pairs, oxygen will have 2 lone pairs, and each fluorine will have 3 lone pairs.

Determine Geometry using VSEPR

Using VSEPR theory, for molecules with four regions of electron density (like COF2 with its double bond and two single bonds), the geometry is predicted to be trigonal planar. However, because the double bond introduces a larger electron cloud, the overall shape will be slightly distorted.

Determine Polarity

Assess the molecular polarity by looking at the arrangement of different electronegative atoms around the carbon. Oxygen is more electronegative than fluorine, resulting in a net dipole moment. The molecule is polar because the presence of the polar bonds and the molecular shape does not cancel out the dipole moments.

Identify Hybridization of COF2

The central carbon atom participates in one double bond and two single bonds, involving three electron groups. This corresponds to sp2 hybridization for the carbon atom.

Make a Sketch showing Orbital Overlap

Sketch the molecule according to valence bond theory showing the sp2 hybrid orbitals of carbon overlapping with one p orbital of oxygen and one p orbital of each fluorine atom. Remember to also depict the unhybridized p orbital of carbon which forms the π bond with oxygen.

Draw Lewis Structure for S2Cl2

For S2Cl2 (disulfur dichloride), sulfur has 6 valence electrons and chlorine has 7 valence electrons. The pi bond will be formed between the two sulfur atoms. Begin by drawing each sulfur atom with a single bond to each other and to a chlorine atom, then fill in the electrons for the sulfur and chlorine atoms to satisfy their octets.

Determine Geometry using VSEPR for S2Cl2

Using VSEPR theory, and considering sulfur as the central atom with two regions of electron density, the molecular geometry of each individual SCl2 unit is bent or V-shaped. However, because there are two such units in S2Cl2, the overall molecule geometry won't be a simple V-shape but will depend on the dihedral angle between the two SCl2 units.

Determine Polarity for S2Cl2

Considering molecular polarity, each SCl2 unit will have a net dipole moment due to the difference in electronegativity between S and Cl and the bent shape of the SCl2. These moments do not cancel out and the molecule is polar.

Identify Hybridization of S2Cl2

The sulfur atoms in S2Cl2 are each bonded to two atoms (the other sulfur and a chlorine), which would imply an sp hybridization. However, lone pairs on S also occupy additional orbitals, leading to sp3 hybridization of the sulfur atoms.

Sketch of S2Cl2 showing Orbital Overlap

Draw the molecule with sp3 hybrid orbitals on sulfur, which overlap with p orbitals on chlorine to form Sigma bonds. The two sulfur atoms will also overlap p orbitals to form the S-S pi bond.

Draw Lewis Structure for SF4

Draw the Lewis structure of SF4 with five regions of electron density (four bonded pairs and one lone pair). Sulfur has 6 valence electrons, and fluorine has 7 valence electrons. The sulfur atom forms single bonds with each of the four fluorine atoms, using all sulfur valence electrons. The sulfur central atom supports a lone pair of electrons.

Determine Geometry using VSEPR for SF4

Using VSEPR theory, SF4 with four bonding pairs and one lone pair of electrons on the sulfur is expected to have a see-saw shaped geometry.

Determine Polarity for SF4

SF4 is polar because the molecule's shape causes the dipoles from the S-F bonds not to cancel each other out, and there is a lone pair on the central atom which adds to the asymmetry of the shape.

Identify Hybridization of SF4

The sulfur atom in SF4 has five regions of electron density, involving sp3d hybridization.

Sketch of SF4 showing Orbital Overlap

Depict SF4 with the central sulfur atom making use of sp3d hybrid orbitals to form sigma bonds with s orbitals of the four fluorine atoms. Also, show the lone pair on sulfur, which occupies an sp3d hybrid orbital.

Key concepts

Lewis Structure

Understanding the Lewis structure of a molecule is the starting point for predicting its properties. It depicts the arrangement of valence electrons around atoms within a molecule, including any bonding pairs between atoms and lone pairs not involved in bonding. The Lewis structure is crucial because it sets the stage for determining other molecular attributes like geometry, polarity, and hybridization. When drawing the Lewis structure, we assign electrons to their respective atoms based on their valence, ensuring that the octet rule is satisfied for each atom. For instance, the exercise provided illustrates drawing Lewis structures for compounds like \textbf{COF}\(_2\), \textbf{S}\(_2\)\textbf{Cl}\(_2\), and \textbf{SF}\(_4\) by allocating electrons to satisfy each element's valence requirements and by creating bonds between atoms.

Molecular Geometry

The molecular geometry of a compound can be predicted using Valence Shell Electron Pair Repulsion (VSEPR) theory. This theory posits that electron groups around the central atom repel each other and arrange themselves as far apart as possible, leading to specific geometric shapes. For example, in a molecule like \textbf{COF}\(_2\) with a central carbon having three groups of electrons, the geometry is trigonal planar. Other configurations, like the see-saw shape of \textbf{SF}\(_4\), emerge from central atoms with different numbers of bonding pairs and lone pairs. The step-by-step solutions detail how the molecular geometry for each given molecule is determined by considering the number of electron groups and the types of electron pairs present.

Polarity of Molecules

The concept of polarity explains how unevenly electrons are distributed in a molecule, resulting in regions of partial positive and negative charges. Polarity arises due to differences in electronegativity between atoms and the molecule's geometric shape. For molecules like \textbf{COF}\(_2\) and \textbf{SF}\(_4\), the dipole moments due to polar bonds don’t cancel out due to the asymmetric shape of the molecule. In these cases, we observe a net dipole moment making the molecule polar. Understanding polarity is essential as it plays a crucial role in the molecule's physical properties, such as boiling and melting points, and chemical properties, including reactivity and solubility.

Hybridization

Hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals that can accommodate the pairing of electrons to form chemical bonds. It's a central idea in valence bond theory that helps us understand the molecular bonding properties. As displayed in the exercise, the hybridization of an atom within a molecule such as \textbf{COF}\(_2\) (sp\(^2\) hybridization) or \textbf{SF}\(_4\) (sp\(^3\)d hybridization) correlates to the number of electron groups surrounding it. Identifying the hybridization helps to picture how the electrons occupy the orbitals and how the atomic orbitals overlap to form molecular bonds.

Valence Bond Theory

Valence bond theory portrays how electron pairs are shared between atoms to form chemical bonds. It emphasizes the significance of orbital overlap for bond formation and provides a detailed approach for visualizing the bonding process. For instance, in the case of \textbf{S}\(_2\)\textbf{Cl}\(_2\), the exercise explains how the sulfur atoms, using sp\(^3\) hybridized orbitals, overlap with p orbitals on chlorine to form sigma bonds. Understanding valence bond theory allows us to conceptualize not only single bonds but also double and triple bonds within complex molecules.

Orbital Overlap

Orbital overlap is a critical concept of valence bond theory where the physical interaction of atomic orbitals from different atoms form covalent bonds. Sigma (\textbf{σ}) bonds show head-to-head overlap, while pi (\textbf{π}) bonds involve side-by-side overlap. In the exercise solutions, students visualize sigma bond formation between the s orbitals of fluorine and the sp\(^3\)d hybrid orbitals of sulfur in \textbf{SF}\(_4\), along with the necessary lone pair positioning. This idea not only reinforces the spatial arrangement of atoms but also relates to the bond strength and molecular geometry.