Question

Predict the structures of (a) \(\left[\mathrm{ICl}_{4}\right]^{-},(\mathrm{b})\left[\mathrm{BrF}_{2}\right]^{+}\) (c) \(\left[\mathrm{ClF}_{4}\right]^{+}\) (d) \(\mathrm{IF}_{7}\) (e) \(\mathrm{I}_{2} \mathrm{Cl}_{6},(\mathrm{f})\left[\mathrm{IF}_{6}\right]^{+},(\mathrm{g}) \mathrm{Br} \mathrm{F}_{5}\)

Short answer

The structures are: (a) Square planar, (b) Bent, (c) Seesaw, (d) Pentagonal bipyramidal, (e) T-shaped, (f) Octahedral, (g) Square pyramidal.

Step-by-step solution

Determine the total valence electrons

For each molecule or ion, count the total number of valence electrons from all the atoms involved, including any added or subtracted electrons due to the charge of the ion.

Calculate steric number

Calculate the steric number, which is the number of atoms bonded to the central atom plus the number of lone pairs on the central atom. This is found by dividing the total valence electrons (from Step 1) by 2, then considering bond formation and leftover electrons.

Predict Geometry Using VSEPR Theory

Use the steric number to predict the molecular geometry using the Valence Shell Electron Pair Repulsion (VSEPR) theory. Correspond the steric number to known geometric shapes (linear, bent, trigonal planar, tetrahedral, etc.).

Assign molecular shapes

Assign molecular shapes to each compound or ion using the geometries from VSEPR theory. Consider factors such as lone pairs causing adjustments from ideal geometries.

Key concepts

Valence Electrons

Valence electrons are the outermost electrons of an atom and play a key role in chemical bonding and molecular structure. These electrons are found in the highest energy level of an atom and are involved in forming chemical bonds with other atoms. We count valence electrons to understand the bonding capacity of an atom, essential when working with VSEPR theory.

To determine the total number of valence electrons, you need to know the group number of each element in the periodic table; this gives the number of valence electrons for the neutral atom. For example, iodine, in group 17, has seven valence electrons. In ions, you must add or subtract electrons based on the charge. For instance, for \([ ext{ICl}_4]^−\), you add an extra electron due to the negative charge, bringing the total to 36 valence electrons for the complete structure.

Steric Number

The steric number is a crucial concept in predicting the shape of molecules using VSEPR theory. It informs us about the spatial arrangement of electron pairs around a central atom. Calculating the steric number involves adding the number of atoms directly bonded to the central atom with the number of lone pairs present.

To determine the steric number efficiently, count the total number of valence electrons and divide by 2 to find the number of electron pairs. Next, identify bonds and lone pairs on the central atom. For example, in \([ ext{ICl}_4]^−\), iodine has a steric number of 6, with 4 bonded atoms and 2 lone pairs, considering the extra electron from the charge. Steric numbers help correlate molecular shapes with known geometric structures.

Molecular Geometry

Molecular geometry is the three-dimensional arrangement of atoms in a molecule as determined by the steric number. Using VSEPR theory, we predict the shape by minimizing electron pair repulsion around a central atom. The goal is to position electron pairs as far apart as possible to reduce repulsive forces.

Molecular geometry is classified into shapes like linear, trigonal planar, tetrahedral, etc. This classification depends on the steric number and presence of lone pairs. For instance, in \([ ext{ICl}_4]^−\), with a steric number of 6 and 2 lone pairs, the geometry is square planar. This adjustment accounts for lone pairs that slightly alter idealized bond angles, highlighting how electron repulsion shapes molecules.

Lone Pairs

Lone pairs refer to pairs of valence electrons that are not involved in chemical bonding. These pairs occupy space around the atom and are vital in determining the shape and bond angles of a molecule.

Lone pairs create electron repulsion, significantly affecting molecular geometry. Unlike bonding pairs, lone pairs take up more space, often compressing bond angles, thus altering ideal molecular shapes. For example, in \([ ext{ICl}_4]^−\), the two lone pairs on iodine heavily influence the final geometry, leading to a square planar structure instead of octahedral, despite the central atom's total coordination number being six. Understanding lone pairs is crucial for accurately predicting deviations from ideal angles and shapes.