Question

These are examples of molecules and ions that do not obey the octet rule. After drawing the Lewis structure, describe the electron-region geometry and the molecular geometry for each. (a) \(\mathrm{ClF}_{2}^{-}\) (b) \(\mathrm{ClF}_{3}\) (c) \(\mathrm{ClF}_{4}^{-}\) (d) \(\mathrm{ClF}_{-}\)

Short answer

(a) Linear, (b) T-shaped, (c) Square planar, (d) Linear.

Step-by-step solution

Draw Lewis Structure for ClF2-

For the ion \( \text{ClF}_2^{-} \), place Cl as the central atom surrounded by the F atoms. The Cl atom has 7 valence electrons and each F atom has 7 valence electrons. Add one extra electron for the negative charge, totaling 22 electrons. Connect Cl and F atoms with single bonds, leaving 18 electrons. Distribute remaining electrons to satisfy the octet for F atoms and then place the remaining as lone pairs on Cl.

Determine Geometry for ClF2-

With three lone pairs and two bonds, \( \text{ClF}_2^{-} \) has five electron regions, resulting in a trigonal bipyramidal electron-region geometry. The actual molecular geometry is linear because the lone pairs push the F atoms farthest apart.

Draw Lewis Structure for ClF3

For \( \text{ClF}_3 \), Cl is the central atom with three F atoms surrounding it. Chlorine has 7 valence electrons and each fluorine has 7, totaling 28 electrons. Form three Cl-F bonds, leaving 20 electrons. Complete the octet for F atoms, then place remaining electrons as lone pairs on Cl.

Determine Geometry for ClF3

\( \text{ClF}_3 \) has five regions including two lone pairs and three bonds, which forms a trigonal bipyramidal electron geometry. The molecular shape is T-shaped due to lone pair repulsion.

Draw Lewis Structure for ClF4-

For \( \text{ClF}_4^{-} \), Cl is the central atom surrounded by four F atoms. Chlorine has 7 valence electrons, each fluorine has 7, plus an extra for the negative charge, totaling 36 electrons. Make four Cl-F bonds, leaving 20 electrons. Complete the octet for F atoms and assign the remaining electrons as lone pairs to Cl.

Determine Geometry for ClF4-

\( \text{ClF}_4^{-} \) has six regions of electron density (two lone pairs and four bonds), leading to an octahedral electron geometry. The molecular geometry is square planar since the two lone pairs are opposite each other.

Draw Lewis Structure for ClF-

For \( \text{ClF}^{-} \), assign Cl as the central atom with F. Cl has 7 valence electrons, F has 7, and there is one additional electron due to the negative charge, making a total of 16 electrons. Form a Cl-F bond, distribute remaining electrons for a complete octet on F, and place remaining electrons as lone pairs on Cl.

Determine Geometry for ClF-

\( \text{ClF}^{-} \) has a total of four electron regions (one bond and three lone pairs), forming a tetrahedral electron-region arrangement. The molecular geometry is linear since lone pairs maximize separation.

Key concepts

Octet Rule

The octet rule is a fundamental concept in chemistry that helps us understand how atoms bond together in molecules. Generally, atoms strive to attain a stable electron configuration, similar to noble gases, which means having eight electrons in their valence shell. However, some molecules, like those involving chlorine and fluorine, occasionally defy this rule. Chlorine, for example, often forms compounds where it has more than eight electrons, such as in

• ClF₂^-
• ClF₃
• ClF₄^-
• ClF^-

This is because chlorine can expand its valence shell to include additional electrons, using the d orbitals available in its electron configuration. Understanding that not all elements strictly follow the octet rule is crucial in predicting the actual molecular structure and properties of these compounds.

Electron-region Geometry

Electron-region geometry, also known as electronic geometry, refers to the spatial arrangement of all electron regions around a central atom in a molecule. These regions consist of both lone pairs and bonds, whether they are single, double, or triple. For instance, in **ClF₂^-**, the geometry is trigonal bipyramidal due to the presence of two bonds and three lone pairs. Another example, **ClF₄^-**, reveals an octahedral electron-region geometry, comprised of four bonds and two lone pairs. Recognizing the electron-region geometry helps in understanding the base shape around the central atom before determining the final molecular geometry, which considers the repulsions between electron pairs.

Molecular Geometry

While the electron-region geometry considers all regions of electron density, molecular geometry strictly focuses on the shape formed by the bonded atoms. Lone pairs exert repulsion on other electrons but are not depicted in the final molecular geometry visuals.
In **ClF₂^-**, although the electron-region geometry is trigonal bipyramidal, the molecular geometry is linear due to the positioning of bonds in response to lone pair repulsion. For **ClF₄^-**, the square planar molecular geometry results from the placement of bonds and lone pairs around the central chlorine atom. By understanding molecular geometry, one can better predict a molecule's polarity and behavior in chemical reactions.

Lone Pairs

Lone pairs are pairs of valence electrons that are not involved in forming bonds within a molecule. They play a significant role in determining the molecular geometry by affecting the spatial arrangement of atoms due to their repulsiveness, which is stronger than that of bonded pairs. For instance, in **ClF₃**, there are two lone pairs on the central chlorine atom. These lone pairs result in more repulsion, causing the bonded fluorine atoms to adopt a T-shaped geometry. Understanding the presence and influence of lone pairs is essential for accurately predicting the shapes and electronic characteristics of compounds. They not only affect geometric configuration but can also impact properties like solubility and reactivity.