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Covalent Bonding an and d Molecular Geomet Geometry ry General Chemistry introduces students to various models and bond theories that predict the likely structure of a molecule and they are: Lewis structures, VSEPR theory, valence bond theory, and molecular orbital theory. This process begins by creating a Lewis structure of a molecule and VSEPR theory is used to determine the electron and molecular geometry of that molecule. A Lewis structure for covalent compounds (two or more nonmetals bonded together) is created using Lewis symbols for each element in the molecule. A Lewis symbol includes the chemical symbol of an element and with its valence electrons drawn as dots surrounding the symbol, as seen in Figure 1 for oxygen.

Figure 1 – Lewis symbol for oxygen

A Lewis structure combines the Lewis symbols for each atom in a molecule to predict its shape. The steps for creating a Lewis structure are: 1. Write the correct skeletal structure of the molecule. The atoms must be arranged in their correct position in the molecule. Some simples rules for arranging atoms are: 1) place the least electronegative atom in a central position, 2) place more electronegative atoms at terminal positions (outside of the molecule), and 3) always place hydrogens in a terminal position because they only form one bond. 2. Calculate the number of valence electrons for the Lewis structure structure. Sum the number of valence electrons from each atom. For example, the total number of valence electrons for water (H2 O) is eight: 1 (from hydrogen) + 1 (from hydrogen) + 6 (from oxygen) = 8. 3. Distribute the electrons among the atoms atoms. Begin by placing two electrons between each pairs of atoms. Distribute the remaining electrons first around terminal atoms, so they have a total of eight (except for hydrogen, which can only have two). Note, the number of electrons associated with an atom include the electrons held in bonds and those surrounding it. Lastly, place remaining electrons around the central atoms. Atoms from period 3 (row 3) and higher may have more than eight electrons. 4. If any atom lacks an octet, form double or triple bonds bonds. Atoms in period 2 (i.e., C, N, O, F, Ne) always require an octet (eight electrons), so if there are not enough electrons available then these atoms can form double or triple bonds to acquire the electrons they need. Electrons on the terminal atoms (except hydrogen) can be moved to form a double bond (two atoms sharing four electrons) and a triple bond (two atoms sharing six electrons), as seen in Figure 2.

Figure 2 - Lewis structure for carbon dioxide. Each carbon and oxygen atom have eight electrons surrounding them.

Once the Lewis structure is created we can begin to make interpretations about the electron and molecular geometry by counting the electron groups around the central atom. An electron group can be a single bond, double bond, triple bond, or lone pair. For example, there are four electron groups around the nitrogen in ammonia (NH3) and three electron groups around the carbon in formaldehyde (COH2), as seen in Figure 3.

Figure 3 - Lewis structure of ammonia (left) and formaldehyde (right).

Valence Shell Electron Pair Repulsion (VSEPR) theory explains that electron groups attached to a central atom are arranged in a way to minimize the repulsion of the negative charges between the electron groups – so they are as far apart as possible. Therefore, electron groups use all of the space around the central atom to spread out. With this understanding, we can predict the geometry of the electron groups around the central atom. For example, the four electron groups around nitrogen in ammonia would be 109.5° apart which would make the electron geometry tetrahedral.

Figure 4 - Electron geometry (left) and molecular geometry (right) for ammonia.

While the electron geometry refers to how the electron groups are oriented, we are more concerned with where the actual atoms are located. Since ammonia has a lone pair on the nitrogen, the electron and molecular geometry are different (Note: when a central atom does not have any lone pairs, then the electron geometry and molecular geometry are the same).

This means the molecular geometry for ammonia is trigonal pyramidal (Figure 4). Table 1 summarizes the electron and molecular geometry as determined by the number of electron groups and lone pairs present. Table 1 - Electron and molecular geometries of molecules

Now that we know how to draw a molecule and predict its shape, let’s take a second to reflect on how different atoms come together to form a molecule. Recall, an atom’s valence electrons are held within s, p, d, and f orbitals and when valence electrons are shared between two atoms a covalent bond is formed. Valence bond theory expands our understanding of bond formation. It states when a half filled orbital from one atom overlaps with a half filled orbital of another atom, the two electrons align with opposite spins which lowers the energy of the atoms resulting in a bond. Let’s explore this concept with carbon, which has four valence electrons (two are in an s orbital and two are in p orbitals). The electrons in the s orbital are paired; however, the two electrons in the p orbitals are unpaired implying carbon can form two bonds. Since, hydrogen has one electron in an s orbital, it can be presumed that two hydrogen atoms can form two bonds with carbon (Figure 5).

Figure 5 - Valence electron configuration for hydrogen and carbon.

However, this explanation is an over-simplification and does not explain why carbon likes to form four bonds. To explain this phenomenon we introduce the concept of atomic orbital hybridization, where atomic orbitals of the central atom combine to form new atomic orbitals. In carbon, one s and three p orbitals combine to form sp3 hybridized orbitals (Figure 6), which are degenerate (have the same energy). Now carbon’s four electrons occupy each sp3 orbital, allowing carbon to form four bonds.

Figure 6 - Hybridization of carbons atomic orbitals.

We can determine the hybridization of the central atom in a familiar way, by making a Lewis structure and counting the number of electron groups around the central atom. The number of electron groups are used to determine the hybridization as seen in the Table 2. Table 2 - Hybridization of the central atom

It is important to note, that occasionally atoms hybridize some of their atomic orbitals leaving the remainder unhybridized. Again let’s refer to carbon in the molecule formaldehyde (COH2). We are aware that carbon has four atomic orbitals, but for this molecule only three atomic orbitals are hybridized (one s and two p) to create sp2 and one p orbital is left unhybridized. The unhybridized p orbital is now available to form a double bond with a p orbital from oxygen (Figure 7). This trend continues if carbon only hybridizes two atomic orbitals (one s and one p) to create two sp orbitals leaving two p orbitals are left unhybridized. In this situation carbon forms a triple bond with another atom. Figure 7- Lewis structure and bonding diagram of formaldehyde (COH2).

Reference List 1. Tro, N. J. Principles of Chemistry: A Molecular Approach, 3rd Edition; Pearson: 2016.

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VSEPR Exercise Exercise Exercise: Draw the Lewis structure for the following molecules in the space provided. Then use the Lewis structure to identify the hybridization of the central, determine the electron and molecular geometry, and list the bond angles.

Example – Molecular Formula: BrF5 Hybridization of the central atom: 𝑠𝑝3 𝑑 2 Electron geometry: Octahedral Molecular geometry: Square pyramidal Bond Angle(s) (list all if there are multiple): 90°

#1 – Molecular Formula: CS2 Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#2 – Molecular Formula: SbOCl Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#3 – Molecular Formula: GeF5− Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#4 – Molecular Formula: SnBr2 Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#5 – Molecular Formula: H3 CCCH Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#6 – Molecular Formula: IO3 F Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#7 – Molecular Formula: KrF2 Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#8 – Molecular Formula: N3− Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#9 – Molecular Formula: XeO2 Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#10 – Molecular Formula: TeF6 Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#11 – Molecular Formula: H3 CCOCH3 Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):

#12 – Molecular Formula: AsCl+ 4 Hybridization of the central atom:

Electron geometry:

Molecular geometry:

Bond Angle(s) (list all if there are multiple):...