Problem 21 - Chemistry and Chemical Reactivity

Question

Draw a Lewis structure for each of the following molecules or ions. Describe the electron-pair geometry and molecular geometry around the central atom. (a) \(\mathrm{ClF}_{2}\) (b) \(\mathrm{ClF}_{3}\) (c) \(\mathrm{CIF}_{4}\) (d) \(\mathrm{CIF}_{5}\)

Step-by-Step Solution

Answer

ClF2: Linear; ClF3: T-shaped; ClF4 (assumed -): Square planar; ClF5: Square pyramidal.

Step 1: Determine Total Valence Electrons for ClF2

For the molecule ClF2, chlorine has 7 valence electrons, and each fluorine also has 7 valence electrons. Thus, the total number of valence electrons is \(7 + 2 \times 7 = 21\) electrons.

Step 2: Draw Lewis Structure for ClF2

Place Cl in the center and single bond it to two F atoms. Distribute the remaining 17 electrons as lone pairs: three pairs on each F atom (6 electrons per F) and five pairs on Cl (10 electrons). This accounts for all 21 valence electrons.

Step 3: Determine Electron-Pair and Molecular Geometry for ClF2

Cl atom has three lone pairs and two bonded pairs (due to the two F atoms), leading to a trigonal bipyramidal electron-pair geometry. However, with three lone pairs, the molecular geometry is linear.

Step 4: Determine Total Valence Electrons for ClF3

For ClF3, Cl has 7 valence electrons, and each F has 7. The total number of valence electrons is \(7 + 3 \times 7 = 28\) electrons.

Step 5: Draw Lewis Structure for ClF3

Place Cl in the center bonded to three F atoms with single bonds. Distribute remaining 19 electrons: three pairs on each F atom (6 electrons per F) and leave two pairs (4 electrons) as lone pairs on Cl.

Step 6: Determine Electron-Pair and Molecular Geometry for ClF3

The Cl atom is surrounded by three bonds and two lone pairs, resulting in a trigonal bipyramidal electron-pair geometry. The removal of lone pairs results in a T-shaped molecular geometry.

Step 7: Determine Total Valence Electrons for ClF4

For ClF4, Cl has 7 valence electrons, and each F atom has 7. The total number of valence electrons is \(7 + 4 \times 7 = 35\) electrons.

Step 8: Draw Lewis Structure for ClF4

Place Cl in the center bonded to four F atoms. Distribute remaining electrons: three pairs on each F atom (6 electrons per F), leaving one electron as a lone electron on Cl.

Step 9: Determine Electron-Pair and Molecular Geometry for ClF4

Consider this as a negatively charged ion (ClF4) since an unusual number of electrons suggests a mistake; possibly meant as ClF4-. Cl has four bonds and in ClF4- would typically have two lone pairs, suggesting a square planar shape.

Step 10: Determine Total Valence Electrons for ClF5

For ClF5, Cl has 7 valence electrons, and each F has 7. The total number of valence electrons is \(7 + 5 \times 7 = 42\) electrons.

Step 11: Draw Lewis Structure for ClF5

Put Cl in the center with five single bonds to F atoms. Distribute remaining electrons: three pairs on each F (6 electrons per F) and one pair of lone electrons on Cl.

Step 12: Determine Electron-Pair and Molecular Geometry for ClF5

With five bonds and one lone pair, the ClF5 forms an octahedral electron pair geometry, but the molecular geometry is square pyramidal due to the lone pair repulsion.

Key Concepts

Valence Electrons

Valence electrons are the outermost electrons in an atom, crucial for understanding chemical bonding. They determine how atoms interact and bind together to form molecules. Consider chlorine (Cl), which has 7 valence electrons, and fluorine (F), sharing the same count of valence electrons.
For molecules, it's essential to total up these valence electrons to draw accurate Lewis structures.

• Molecules mentioned include \( \text{ClF}_2 \), \( \text{ClF}_3 \), \( \text{ClF}_4^- \), and \( \text{ClF}_5 \).
• Each fluorine atom contributes 7 electrons.
• Chlorine starts with 7 electrons as well.

To compute total valence electrons:

• For \( \text{ClF}_2 \), calculate \((7 + 2 \times 7) = 21 \) electrons.
• For \( \text{ClF}_3 \, \text{(7 + 3 \times 7) = 28} \) electrons.
• For \( ext{ClF}_4^- \, \text{35 electrons,}\) accounting for an extra electron due to negative charge.
• Lastly, \( ext{ClF}_5 \) has \(42 \) electrons with five bonds and the remaining as lone pairs.

Understanding valence electrons allows us to accurately draw Lewis structures and predict molecular behavior.

Molecular Geometry

Molecular geometry describes the three-dimensional shape formed by a molecule's atoms, focusing only on the positions of the nuclei and bonded atoms. It doesn't consider the lone pairs, but these do affect the resultant shape.
For example, in the structure of \(\text{ClF}_2\):

• Despite having a trigonal bipyramidal electron geometry, \(\text{ClF}_2\) manifests as a linear shape due to equivalent dispersion of electron bonds and repulsions by the lone pairs.

Similarly, take \(\text{ClF}_3\):

• The electron-pair geometry is trigonal bipyramidal. However, the arrangement of three fluorine atoms around central Cl with two lone pairs results in a T-shaped structure.

Further, for ions like \(\text{ClF}_4^-\):

• Though the electron geometry indicates a square planar, influenced by repulsion balance.

And for \(\text{ClF}_5\):

• We expect an octahedral shape initially; yet, with lone pairs, it shifts to square pyramidal.

Understanding these geometries involves knowing how lone pairs alter molecule shape based on VSEPR theory, affecting intermolecular interactions and reactivity.

Electron-Pair Geometry

Electron-pair geometry considers not only the atoms but also the electron pairs, including lone pairs, forming around a central atom. This provides insight into the spatial arrangement that can influence molecular geometry.
When evaluating \(\text{ClF}_2\):

• The geometry includes three lone pairs and two bonded pairs indicating a trigonal bipyramidal shape.

For \(\text{ClF}_3\):

• We see three bonds and two lone pairs, maintaining the trigonal bipyramidal electron geometry.

Similarly, in \(\text{ClF}_4^-\):

• Four bonds and two lone pairs result in the square planar arrangement.

For \(\text{ClF}_5\):

• An octahedral configuration, results initially; however, the lone pairs shift arrangements to square pyramidal.

Studying electron-pair geometry allows for an early indicator of potential molecular shape modifications. It explains how electron repulsions influence molecular form, improving predictions of molecular interactions.