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Chapter 9: Problem 13

How does a trigonal pyramid differ from a tetrahedron so far as molecular geometry is concerned?

Short Answer

Expert verified
A trigonal pyramidal molecular geometry has one central atom with three surrounding atoms at the corners of a pyramid and a lone pair of electrons on the central atom. The bond angles are typically around 107.5 degrees, and the molecule is polar due to the asymmetric electron distribution. In contrast, a tetrahedral molecular geometry has one central atom with four surrounding atoms at the vertices of a regular tetrahedron. The bond angles are always 109.5 degrees, and the molecule is often nonpolar due to the absence of lone pairs on the central atom and symmetrical distribution of outer atoms. Examples include ammonia (NH鈧�) for trigonal pyramidal and methane (CH鈧�) for tetrahedral.

Step by step solution

01

1. Defining Trigonal Pyramid and Tetrahedron

A trigonal pyramid is a molecular geometry with one atom at the center and three surrounding atoms (3 outer atoms) at the corners of a pyramid, and it also has a lone pair of electrons at the central atom. This shape is also known as a trigonal pyramidal. A tetrahedron, on the other hand, is a molecular geometry with one central atom and four surrounding atoms (4 outer atoms) at the vertices of a regular tetrahedron. This shape is also referred to as a tetrahedral.
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2. Differences in Bond Angles

In a trigonal pyramidal geometry, the bond angles between the surrounding atoms are typically around 107.5 degrees because of the repulsion caused by the lone pair of electrons on the central atom. The ideal angle would be equal if there were no lone pair present, which is 109.5 degrees. In a tetrahedral geometry, the bond angles are always 109.5 degrees because there are no lone pairs on the central atom to cause additional repulsion.
03

3. Differences in Symmetry

A trigonal pyramidal molecule is polar because it has an asymmetric electron distribution with the presence of lone pairs. This means that there is a net dipole moment in the molecule, causing them to have a polar nature. In contrast, a tetrahedral molecule is often nonpolar because there are no lone pairs on the central atom, and the distribution of outer atoms is symmetrical. Therefore, in a pure tetrahedral molecule, the net dipole moment is zero, making it nonpolar.
04

4. Examples of Trigonal Pyramidal and Tetrahedral Molecules

An example of a trigonal pyramidal molecule is ammonia (NH鈧�), where nitrogen is the central atom bonded with three hydrogen atoms, and there is a lone pair of electrons on the nitrogen. An example of a tetrahedral molecule is methane (CH鈧�), where the central atom is carbon bonded to four hydrogen atoms, and there are no lone pairs present on the central atom.
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Conclusion

In summary, although trigonal pyramidal and tetrahedral molecular geometries might seem similar, they have significant differences. The primary differences arise from the presence of a lone pair on the central atom in a trigonal pyramidal, which causes deviations in bond angles, polarity, and symmetric molecular structure compared to a tetrahedral molecule, which has no lone pairs on the central atom.

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Most popular questions from this chapter

Consider the Lewis structure for glycine, the simplest amino acid: (a) What are the approximate bond angles about each of the two carbon atoms, and what are the hybridizations of the orbitals on each of them? (b) What are the hybridizations of the orbitals on the two oxygens and the nitrogen atom, and what are the approximate bond angles at the nitrogen? (c) What is the total number of \(\sigma\) bonds in the entire molecule, and what is the total number of \(\pi\) bonds?

Explain the following: (a) The peroxide ion, \(\mathrm{O}_{2}^{2-}\), has a longer bond length than the superoxide ion, \(\mathrm{O}_{2}^{-}\). (b) The magnetic properties of \(\mathrm{B}_{2}\) are consistent with the \(\pi_{2 p}\) MOs being lower in energy than the \(\sigma_{2 p}\) MO. (c) The \(\mathrm{O}_{2}^{2+}\) ion has a stronger O \(-\) O bond than \(\mathrm{O}_{2}\) itself.

(a) Explain why \(\mathrm{BrF}_{4}^{-}\) is square planar, whereas \(\mathrm{BF}_{4}^{-}\) is tetrahedral. (b) How would you expect the \(\mathrm{H}-\mathrm{X}-\mathrm{H}\) bond angle to vary in the series \(\mathrm{H}_{2} \mathrm{O}, \mathrm{H}_{2} \mathrm{~S}, \mathrm{H}_{2} \mathrm{Se}\) ? Explain. (Hint: The size of an electron pair domain depends in part on the electronegativity of the central atom.)

The \(\mathrm{O}-\mathrm{H}\) bond lengths in the water molecule \(\left(\mathrm{H}_{2} \mathrm{O}\right)\) are \(0.96 \AA\), and the \(\mathrm{H}-\mathrm{O}-\mathrm{H}\) angle is \(104.5^{\circ} .\) The dipole moment of the water molecule is \(1.85 \mathrm{D} .\) (a) In what directions do the bond dipoles of the \(\mathrm{O}-\mathrm{H}\) bonds point? In what direction does the dipole moment vector of the water molecule point? (b) Calculate the magnitude of the bond dipole of the \(\mathrm{O}-\mathrm{H}\) bonds. (Note: You will need to use vector addition to do this.) (c) Compare your answer from part (b) to the dipole moments of the hydrogen halides (Table 8.3). Is your answer in accord with the relative electronegativity of oxygen?

(a) Methane \(\left(\mathrm{CH}_{4}\right)\) and the perchlorate ion \(\left(\mathrm{ClO}_{4}^{-}\right)\) are both described as tetrahedral. What does this indicate about their bond angles? (b) The \(\mathrm{NH}_{3}\) molecule is trigonal pyramidal, while \(\mathrm{BF}_{3}\) is trigonal planar. Which of these molecules is flat?
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