Question

Draw a plausible Lewis structure for the following series of molecules and ions: (a) \(\mathrm{ClF}_{2}^{-} ;\) (b) \(\mathrm{ClF}_{3}\); (c) \(\mathrm{ClF}_{4}^{-} ;\) (d) ClF \(_{5} .\) Describe the electron group geometry and molecular structure of these species.

Short answer

The Lewis Structures of the molecules are derived from the total valence electrons. For (a) \(\mathrm{ClF}_{2}^{-}\), (b) \(\mathrm{ClF}_{3}\), (c) \(\mathrm{ClF}_{4}^{-}\) and (d) \(\mathrm{ClF}_{5}\), the electron group geometries are Trigonal Planar, Trigonal Bipyramidal, Trigonal Bipyramidal, and Octahedral respectively. The molecular structures are T-shape, T-shape, Seesaw and Octahedral respectively.

Step-by-step solution

(a): Lewis Structure for \(\mathrm{ClF}_{2}^{-}\)

First, count total valence electrons present: Chlorine (7 electrons) + 2*Fluorine (7 electrons each) + 1 electron (for the -1 charge) = 22 electrons. Draw a skeleton wherein a Chlorine atom is central and is bonded with two Fluorines. Subtract the electrons used for bonding from total electrons: 22-4 =18. Distribute the remaining electrons on atoms, beginning with outer atoms. Each outer Fluorine atom receives 6 electrons and Chlorine gets the remaining 6 electrons. Now, verify that each atom follows Octet rule (except Hydrogen). The structure following the octet rule is the correct Lewis Structure.

(b): Lewis Structure for \(\mathrm{ClF}_{3}\)

First, count total valence electrons present: Chlorine (7 electrons) + 3*Fluorine (7 electrons each) = 28 electrons. Draw a skeleton where Chlorine atom is central and bonded with three Fluorines. Subtract bonding electrons from total: 28-6 =22. Distribute the remaining electrons on atoms. Each outer Fluorine atom receives 6 electrons and Chlorine gets the remaining 4 electrons. Verify that each atom follows Octet rule.

(c): Lewis Structure for \(\mathrm{ClF}_{4}^{-}\)

First, count total valence electrons present: Chlorine (7 electrons) + 4*Fluorine (7 electrons each) +1 electron (for -1 charge) = 36 electrons. Draw a skeleton with Chlorine atom as central and bonded with four Fluorines. Subtract bonding electrons from total: 36-8 = 28. Distribute the remaining electrons on atoms. Each outer Fluorine atom receives 6 electrons and Chlorine gets the remaining 4 electrons.

(d): Lewis Structure for ClF \(_{5}\)

The total valence electrons present are Chlorine (7 electrons) + 5*Fluorine (7 electrons each) = 42 electrons. Draw a skeleton where Chlorine atom is central and bonded with five Fluorines. Subtract bonding electrons: 42-10 = 32. Distribute the remaining electrons on atoms. Each outer Fluorine atom receives 6 electrons.

Electron Group Geometry and Molecular Structure

Electron group geometry considers both bonding and non-bonding electron pairs, while molecular structure considers only bonding pairs. (a) For \(\mathrm{ClF}_{2}^{-}\), there are 3 electron groups (2 bonding and 1 lone pair) around central Chlorine. So its group geometry is T-shape and electron geometry is Trigonal Planar. (b) For \(\mathrm{ClF}_{3}\), there are 5 electron groups (3 bonding and 2 lone pairs). Hence, its group geometry is T-shape and electron geometry is Trigonal Bipyramidal. (c) For \(\mathrm{ClF}_{4}^{-}\), there are 5 electron groups (4 bonding and 1 lone pair). Hence, its group geometry is Seesaw and electron geometry is Trigonal Bipyramidal. (d) For \(\mathrm{ClF}_{5}\), there are 6 electron groups (all bonding), its group geometry and molecular structure are Octahedral.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom, which are involved in forming chemical bonds. The number of valence electrons an element possesses can greatly influence the chemical reactivity and bonding capacity of that element. When creating Lewis structures, it's crucial to accurately determine the valence electrons.

• Chlorine (Cl) has 7 valence electrons.
• Fluorine (F) also has 7 valence electrons.
• For anions like \(\mathrm{ClF}_{2}^{-}\), an extra electron is added to account for the negative charge, increasing the total valence electrons by 1.

To illustrate, in the case of \(\mathrm{ClF}_{3}\), we start with 7 electrons from chlorine and 21 from the three fluorines, totaling 28 electrons. Correctly distributing these electrons ensures each atom follows the octet rule, where possible, ensuring stability in the resulting structure.

Electron Geometry

Electron geometry describes the spatial arrangement of all electron groups surrounding a central atom, including lone pairs. This geometry is determined by the VSEPR (Valence Shell Electron Pair Repulsion) theory, which predicts that electron pairs will arrange themselves as far apart as possible to minimize repulsion.

• For \(\mathrm{ClF}_{2}^{-}\), we assess three electron groups: two bonding pairs and one lone pair. This configuration results in a T-shaped electron geometry.
• With \(\mathrm{ClF}_{3}\), there are five electron groups: three bonding pairs and two lone pairs. This arrangement leads to a trigonal bipyramidal electron geometry.
• In \(\mathrm{ClF}_{4}^{-}\), with four bonding and one lone pair, its electron geometry is described as seesaw, again rooted in a trigonal bipyramidal shape.
• For \(\mathrm{ClF}_{5}\), six bonding pairs around the central atom give it an octahedral electron geometry.

Understanding electron geometry helps us predict and explain the shape and angles between bonds in molecules.

Molecular Structure

Molecular structure refers to the actual 3D arrangement of atoms in a molecule. Unlike electron geometry, molecular structure considers only the arrangement of bonded atoms, ignoring lone pairs, which can alter the shape perception.
In the case of \(\mathrm{ClF}_{2}^{-}\), even though its electron geometry is trigonal planar, the presence of a lone pair makes its molecular structure T-shaped. Similarly, \(\mathrm{ClF}_{3}\), despite having a trigonal bipyramidal electron geometry, takes on a T-shaped molecular structure due to its lone pairs.
Seesaw geometry in compounds like \(\mathrm{ClF}_{4}^{-}\) emerges when one of the positions in a trigonal bipyramidal shape is occupied by a lone pair. Lastly, \(\mathrm{ClF}_{5}\) boasts a perfectly octahedral structure, with no lone pairs to alter its symmetric bond arrangement.
Recognizing molecular structure is important for determining properties like polarity and reactivity of molecules. Each unique arrangement contributes significantly to how a molecule will interact chemically and physically in various environments.