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Chapter 10: Problem 59

Which of these species is not likely to have a tetrahedral shape (a) \(\operatorname{SiBr}_{4}\) (b) \(\mathrm{NF}_{4}^{+},(\mathrm {c}) \mathrm{SF}_{4}\) (d) \(\mathrm{BeCl}_{4}^{2-}\) (e) \(\mathrm{BF}_{4}^{-},\) (f) \(\mathrm{AlCl}_{4}^{-}\)

Short Answer

Expert verified
The species \( \mathrm{NF}_{4}^{+}\) and \( \mathrm{SF}_{4}\) are not likely to have a tetrahedral shape.

Step by step solution

01

- Determine the Electron Distribution

To solve this, first determine the electron distribution around the central atom for each option. Using the octet rule, recall that atoms will gain, lose, or share electrons to achieve a stable electronic configuration with 8 electrons in their outermost (valence) shell.
02

- Apply VSEPR Theory

Next, apply the VSEPR theory to predict the shape of the molecule based on the number of valence electrons and their distribution. In a tetrahedral shape, the central atom is surrounded by 4 other atoms or groups of atoms, all of which are spaced as far apart as possible because of electron pair repellence.
03

- Evaluate Each Option

Evaluate the molecules: (a) \( \operatorname{SiBr}_{4}\) - silicon atom has four bromine atoms connected, tetrahedral; (b) \( \mathrm{NF}_{4}^{+}\) - nitrogen atom has four fluorine atoms and a lone pair of e-, not tetrahedral; (c) \( \mathrm{SF}_{4}\) - sulfur atom has four fluorine atoms and a lone pair of e-, not tetrahedral; (d) \( \mathrm{BeCl}_{4}^{2-}\) - beryllium atom equipped with four chloride ions, tetrahedral; (e) \( \mathrm{BF}_{4}^{-}\) - boron atom connected with four fluorine atoms, tetrahedral; (f) \( \mathrm{AlCl}_{4}^{-}\) - aluminum atom surrounded by four chlorine atoms, tetrahedral.
04

- Identify the Non-Tetrahedral Shape

From the evaluation in step 3, it can be seen that \( \mathrm{NF}_{4}^{+}\) and \( \mathrm{SF}_{4}\) do not have a tetrahedral shape due to the presence of a lone pair of electrons on the central atom. Therefore, these are the species that do not have a tetrahedral shape.

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Most popular questions from this chapter

Describe the geometry around each of the three central atoms in the \(\mathrm{CH}_{3} \mathrm{COOH}\) molecule.

What hybrid orbitals are used by nitrogen atoms in these species? (a) \(\mathrm{NH}_{3},\) (b) \(\mathrm{H}_{2} \mathrm{~N}-\mathrm{NH}_{2}\) (c) \(\mathrm{NO}_{3}^{-}\)

Predict the bond angles for these molecules: (a) \(\mathrm{BeCl}_{2},\) (b) \(\mathrm{BCl}_{3},\) (c) \(\mathrm{CCl}_{4},\) (d) \(\mathrm{CH}_{3} \mathrm{Cl}\) (e) \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}\) (arrangement of atoms: \(\mathrm{ClHgHgCl}\) ), (f) \(\mathrm{SnCl}_{2}\) (g) \(\mathrm{H}_{2} \mathrm{O}_{2},\) (h) \(\mathrm{SnH}_{4}\).

Draw Lewis structures and give the other information requested for the following molecules: (a) \(\mathrm{BF}_{3}\). Shape: planar or nonplanar? (b) \(\mathrm{ClO}_{3}^{-}\). Shape: planar or nonplanar? (c) \(\mathrm{H}_{2} \mathrm{O}\). Show the direction of the resultant dipole moment. (d) \(\mathrm{OF}_{2}\). Polar or nonpolar molecule? (e) \(\mathrm{SeO}_{2}\). Estimate the OSeO bond angle.

Which of these species has a longer bond, \(\mathrm{B}_{2}\) or \(\mathrm{B}_{2}^{+} ?\) Explain in terms of molecular orbital theory.
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