Question

Predict the molecular structure for each of the following. (See Exercises 115 and 116.) a. ${\mathrm{BrFI}}_2$ b. ${\mathrm{XeO}}_2{\mathrm{F}}_2$ c. ${\mathrm{TeF}}_2{\mathrm{Cl}}_3^{-}$ For each formula there are at least two different structures that can be drawn using the same central atom. Draw all possible structures for each formula.

Short answer

In summary, the molecular structures for each compound are: a. BrFI2: Two structures with trigonal bipyramidal electron domain geometry, differing in the placement of F and I atoms in axial or equatorial positions. b. XeO2F2: Two structures with tetrahedral electron domain geometry, differing in the placement of O and F atoms or O, F, and a lone pair in tetrahedral positions. c. TeF2Cl3-: Two structures with octahedral electron domain geometry, differing in the placement of Cl atoms and F atoms or Cl atoms and a lone pair in axial or equatorial positions.

Step-by-step solution

1. Determine Lewis Structure for BrFI2

First, we need to determine the number of valence electrons for each element in the compound. For Br (Bromine), there are 7 valence electrons; for F (Fluorine), there are also 7 valence electrons; for I (Iodine), there are 7 valence electrons. So the total valence electrons are 7 + 7 + 7*2 = 28. To draw the Lewis structure, put Br in the center and surround it with three F atoms and I atoms. Now distribute the valence electrons to achieve an octet for each atom while minimizing formal charges.

2. Determine Electron Domain Geometry for BrFI2

Using the Lewis structure, we have 3 bonding electron pairs and 2 lone pairs of electrons around the central Br atom. Thus, the electron domain geometry is trigonal bipyramidal, with an electron domain count of 5.

3. Determine Molecular Geometry for BrFI2

In the trigonal bipyramidal electron domain geometry, two 3-electron domain atoms can occupy axial positions and other three can occupy equatorial positions. Take two arrangements for BrF2I: - In the first one, put two F atoms in axial positions, and an I atom in an equatorial position. - In the second one, put an F atom and an I atom in axial positions, and an I atom in an equatorial position. Note that these structures are only different if F and I were not the same type of atoms.

4. Repeat Steps 1-3 for XeO2F2

- Total valence electrons: 8 (Xe) + 6 (O) * 2 + 7(F) * 2 = 34 - Electron domain geometry: Tetrahedral (4 electron domains) - Molecular geometry: Two arrangements: 1. Xe in the center with two O atoms and two F atoms in tetrahedral positions. 2. Xe in the center with one O atom, one F atom, and one lone pair in tetrahedral positions.

5. Repeat Steps 1-3 for TeF2Cl3-

- Total valence electrons: 6 (Te) + 7(F) * 2 + 7(Cl) * 3 + 1 (extra electron due to -1 charge) = 42 - Electron Domain Geometry: Octahedral (6 electron domains) - Molecular Geometry: Two arrangements: 1. Te in the center with Cl atoms in axial positions, F atoms and one lone pair in equatorial positions. 2. Te in the center with Cl atoms and one lone pair in axial positions, F atoms in equatorial positions. In conclusion, the molecular structures for each of the following are: a. BrFI2: Two structures with trigonal bipyramidal electron domain geometry (described in step 3). b. XeO2F2: Two structures with tetrahedral electron domain geometry (described in step 4). c. TeF2Cl3-: Two structures with octahedral electron domain geometry (described in step 5).

Key concepts

Lewis Structure

The Lewis structure is a useful way to represent molecules showing how atoms are bonded and how electrons are distributed in the molecule. Start by calculating the total number of valence electrons available. Each element's group number in the periodic table usually indicates how many valence electrons they have. For instance, Bromine, Fluorine, and Iodine each have 7 valence electrons. Add these up for each atom in the formula to get the total number of electrons needed.

Next, choose a central atom, typically the one that can make the most bonds or is less electronegative. Position the other atoms around it and use lines to represent pairs of bonding electrons between them. Finally, use the remaining valence electrons to complete the octet rule (8 electrons) for each atom, or duet for hydrogen.

Keep in mind: some elements can hold more than 8 electrons. When distributing electrons, try to minimize formal charges for more stable structures.

Electron Domain Geometry

Electron domain geometry considers the spatial arrangement of all pairs of electrons, both bonding and non-bonding, around the central atom. This is derived from the logical grouping around a central atom in the molecule, often referred to as "electron domains". Each bonded atom and each pair of lone pair electrons around the central atom form their own domain.

For example, in BrFI extsubscript{2}, with three bonding pairs and two lone pairs on Bromine, there are five domains in total. These domains arrange themselves as far apart as possible, resulting in a trigonal bipyramidal geometry. This minimization of repulsion is based on the VSEPR (Valence Shell Electron Pair Repulsion) theory.

• 2 domains: linear
• 3 domains: trigonal planar
• 4 domains: tetrahedral
• 5 domains: trigonal bipyramidal
• 6 domains: octahedral

Remember, electron domain geometry does not always reflect the actual "shape" that appears in chemistry, which is instead described by molecular geometry.

Valence Electrons

Valence electrons are the outermost electrons of an atom that participate in chemical bonding. They are crucial in determining how atoms will bond with each other to form molecules. These electrons are responsible for the chemical properties of the elements, as they are the electrons involved in forming bonds.

To determine the number of valence electrons, look at an element's position on the periodic table. For example, elements in Group 1 have one valence electron, Group 2 elements have two, and so on, up to the Group 18 noble gases which have eight valence electrons (except Helium, which has two).

When forming molecules, atoms will generally share, gain, or lose electrons to reach a more stable electronic configuration, often achieving a so-called "octet" in their outer shells. This stability can explain why, for instance, Na bonds with Cl to form NaCl, complete with full valence shells.

Bonding Electrons

Bonding electrons are the shared electrons in covalent bonds between atoms within a molecule. When two atoms share a pair of electrons, they form a covalent bond, symbolized in Lewis structures as lines between atoms.

The nature of bonding electrons is significant because they determine the strength, length, and angle of the bonds between atoms. They are centrally involved in defining the molecular geometry alongside lone pairs.

In BrFI extsubscript{2}, among the 28 available electrons, the three bonds (Br-F and Br-I) are formed by sharing pairs of electrons. Each bond is formed by two bonding electrons. The way these electrons are shared can be even, as in nonpolar bonds, or uneven among different elements, resulting in polar covalent bonds.

Understanding bonding electrons and their roles help predict not only the stability of molecules but also their reactivity, solubility, and other physical properties.