Problem 124 - Chemistry: An Atoms-Focused Approach

Question

Fluoroaluminate anions \(\mathrm{AIF}_{4}^{-}\) and \(\mathrm{AIF}_{6}^{3-}\) have been known for over a century, but the structure of the pentafluoroaluminate ion, AlFs \(_{5}^{2-},\) was not determined until 2003. Draw the Lewis structures for AlF \(_{3},\) AlF \(_{4}^{-},\) AlF \(_{5}^{2-}\), and AlF \(_{6}^{3-} .\) Determine the molecular geometry of each molecule or ion. Describe the bonding in \(\mathrm{AlF}_{3}, \mathrm{AlF}_{4}^{-}\) \(\mathrm{AlF}_{5}^{2-},\) and \(\mathrm{AlF}_{6}^{3-}\) using valence bond theory.

Step-by-Step Solution

Answer

Question: Describe the molecular geometry and bonding of AlF3, AlF4⁻, AlF5²⁻, and AlF6³⁻ using Lewis structures and valence bond theory. Answer: AlF3 has a trigonal planar geometry and forms 3 sigma bonds with Fluorine atoms using its 3s and 3p orbitals. AlF4⁻ has a tetrahedral geometry and forms 4 sigma bonds with Fluorine atoms using sp³d hybrid orbitals. AlF5²⁻ has a trigonal bipyramidal geometry, and forms 5 sigma bonds with Fluorine atoms using sp³d² hybrid orbitals. AlF6³⁻ has an octahedral geometry and forms 6 sigma bonds with Fluorine atoms using its sp³d² hybrid orbitals.

Step 1: Draw the Lewis structures of each molecule/ion

To draw the Lewis structures, first, we need to know the number of valence electrons for each element. Aluminum (Al) has 3 valence electrons, and Fluorine (F) has 7 valence electrons. 1. AlF3 In AlF3, the total number of valence electrons is (3+3×7) = 24. Aluminum is the central atom surrounded by three Fluorine atoms. Thus, the structure is as follows: Al is surrounded by 3 single bonds to F atoms, and each F atom has 3 lone pairs. 2. AlF4⁻ In AlF4⁻, the total number of valence electrons is (3+4×7)+1(charge) = 32. Aluminum is the central atom surrounded by four Fluorine atoms. Thus, the structure is as follows: Al is surrounded by 4 single bonds to F atoms, and each F atom has 3 lone pairs. 3. AlF5²- In AlF5²-, the total number of valence electrons is (3+5×7)+2(charge) = 40. Aluminum is the central atom surrounded by five Fluorine atoms. Thus, the structure is as follows: Al is surrounded by 5 single bonds to F atoms, and each F atom has 3 lone pairs. 4. AlF6³- In AlF6³-, the total number of valence electrons is (3+6×7)+3(charge) = 48. Aluminum is the central atom surrounded by six Fluorine atoms. Thus, the structure is as follows: Al is surrounded by 6 single bonds to F atoms, and each F atom has 3 lone pairs.

Step 2: Determine the molecular geometry

Once we have the Lewis structures, we can now determine the molecular geometries using the VSEPR (Valence Shell Electron Pair Repulsion) theory. 1. AlF3: Aluminum is surrounded by 3 single bonds and no lone pairs, so it has a trigonal planar geometry. 2. AlF4⁻: Aluminum is surrounded by 4 single bonds and no lone pairs, so it has a tetrahedral geometry. 3. AlF5²-: Aluminum is surrounded by 5 single bonds and no lone pairs, so it has a trigonal bipyramidal geometry. 4. AlF6³-: Aluminum is surrounded by 6 single bonds and no lone pairs, so it has an octahedral geometry.

Step 3: Describe the bonding using valence bond theory

Valence bond theory can be used to explain the bonding in these ions. 1. AlF3: Aluminum uses its 3 valence orbitals (one 3s and two 3p orbitals) to form 3 sigma bonds with three Fluorine atoms. The remaining three lone pairs of electrons occupy other orbitals of the Fluorine atoms. 2. AlF4⁻: Aluminum promotes one of its 3s electrons to the empty 3d orbital and hybridizes its one 3s, three 3p, and one 3d orbitals to form five sp³d hybrid orbitals. It then uses four of these orbitals to form sigma bonds with the four Fluorine atoms. The remaining lone pairs of electrons occupy other orbitals of the Fluorine atoms. 3. AlF5²-: Aluminum promotes one of its 3s electrons to a higher energy level and hybridizes its one 3s, three 3p, and two 3d orbitals to form six sp³d² hybrid orbitals. It then uses five of these orbitals to form sigma bonds with the five Fluorine atoms. The remaining lone pairs of electrons occupy other orbitals of the Fluorine atoms. 4. AlF6³-: Aluminum forms six sigma bonds with six Fluorine atoms, using its sp³d² hybrid orbitals formed after promoting one electron from the 3s orbital to a higher energy level and hybridizing the 3s, 3p, and 3d orbitals. The lone pairs of electrons in the Fluorine atoms occupy other orbitals.

Key Concepts

Lewis Structures

Lewis structures are diagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule. They play a vital role in predicting the geometry and properties of molecules.
To draw Lewis structures for the fluoroaluminate ions, it’s important to know the valence electrons.

• AlF extsubscript{3}: Aluminum has three valence electrons, and each Fluorine has seven. This gives a total of 24 electrons. Aluminum is surrounded by three Fluorine atoms with single bonds, and each Fluorine has three lone pairs.
• AlF extsubscript{4} extsuperscript{-}: With an added electron due to the charge, there are 32 electrons. Aluminum is central with four Fluorine atoms bonded, each with three lone pairs.
• AlF extsubscript{5} extsuperscript{2-}: Including two more electrons due to its charge, the total is 40 electrons. Five single bonds connect the central Aluminum to Fluorines, each still having three lone pairs.
• AlF extsubscript{6} extsuperscript{3-}: It carries one more electron for a total of 48 electrons, forming six single bonds with six Fluorines each having three lone pairs.

These structures are foundational in understanding the molecular shapes and bonding detailed in other theories like VSEPR and Valence Bond Theory.

Molecular Geometry

Molecular geometry describes the three-dimensional arrangement of atoms within a molecule. It helps us predict molecular shapes, angles, and polarities.
For AlF extsubscript{n} ions, the VSEPR theory aids in evaluating molecular geometries:

• AlF extsubscript{3}: AlF extsubscript{3} has a trigonal planar shape since its central Aluminum atom forms three bonds with no lone pairs. The shape is flat, with bond angles of 120°.
• AlF extsubscript{4} extsuperscript{-}: With four bonds around Aluminum, this ion exhibits a tetrahedral geometry, which is characterized by bond angles of about 109.5°.
• AlF extsubscript{5} extsuperscript{2-}: Five bonds lead to a trigonal bipyramidal structure, where angles are split into 90° and 120°. This structure features a clearer distinction between axial and equatorial positions.
• AlF extsubscript{6} extsuperscript{3-}: Here, the octahedral geometry results from six equal bonds around Aluminum, providing 90° angles between bonds.

Understanding the geometry aids in predicting reaction possibilities and the molecule's reactivity.

Aluminum Compounds

Aluminum compounds, especially fluoroaluminates like AlF extsubscript{4} extsuperscript{-}, AlF extsubscript{5} extsuperscript{2-}, and AlF extsubscript{6} extsuperscript{3-}, are known for their unique properties and applications. Aluminum, a group 13 element, often exhibits a +3 oxidation state in its compounds, making it adept at forming strong bonds with halogens such as fluorine.
Fluoride ions, known for their high electronegativity, establish robust interactions with Aluminum, offering stability to these complexes. As these fluoroaluminates coordinate with their central Aluminum atom through sigma bonds, they exhibit distinctive shapes and properties.
These aluminum compounds play a significant role in industrial applications, like catalysis, glass production, and the nuclear industry, where they’re often used for their ability to alter physicochemical properties effectively. The diversity in structure and stability of different Al-fluoro complexes highlights the versatility of aluminum in chemical engineering and materials science.

VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory is a crucial framework for understanding the 3D formation of molecules. According to VSEPR, electron pairs around a central atom tend to arrange themselves to minimize repulsion, thereby determining the molecular geometry.
When applied to the fluoroaluminate ions:

• AlF extsubscript{3}: The three bond pairs form a trigonal planar shape.
• AlF extsubscript{4} extsuperscript{-}: With four bond pairs and no lone pairs, the structure is tetrahedral.
• AlF extsubscript{5} extsuperscript{2-}: The five bond pairs lead to a trigonal bipyramidal shape, with bond pairs being both equatorial and axial.
• AlF extsubscript{6} extsuperscript{3-}: Six bond pairs adopt an octahedral arrangement.

VSEPR theory is essential, allowing chemists to predict molecules' shapes from their formulae, anticipate their interactions, and visualize their spatial configuration. This understanding is particularly crucial in fields like molecular chemistry, where shape greatly influences chemical behavior.