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Chapter 9: Problem 49

Consider the molecule \(\mathrm{BF}_{3}\). (a) What is the electron configuration of an isolated \(\mathrm{B}\) atom? (b) What is the electron configuration of an isolated F atom? (c) What hybrid orbitals should be constructed on the B atom to make the B-F bonds in \(\mathrm{BF}_{3}\) ? (d) What valence orbitals, if any, remain unhybridized on the \(\mathrm{B}\) atom in \(\mathrm{BF}_{3} ?\)

Short Answer

Expert verified
The electron configurations of an isolated B atom and an isolated F atom are \(1s^{2}2s^{2}2p^{1}\) and \(1s^{2}2s^{2}2p^{5}\), respectively. In the \(\mathrm{BF}_{3}\) molecule, the B atom undergoes sp2 hybridization to form three sp2 orbitals for bonding with three F atoms, leaving the 2p_z orbital unhybridized.

Step by step solution

01

Determine the electron configuration of an isolated B atom

To determine the electron configuration of an isolated Boron (B) atom, we need to consider the number of electrons and the order in which they fill the atomic orbitals. Boron has an atomic number of 5, which means it has 5 electrons. The order in which orbitals fill is 1s, 2s, 2p, 3s, 3p, and so on. Following this order, the electron configuration for Boron is: \(1s^{2}2s^{2}2p^{1}\).
02

Determine the electron configuration of an isolated F atom

Now, let's determine the electron configuration of an isolated Fluorine (F) atom. Fluorine has an atomic number of 9, which means it has 9 electrons. Following the same orbital filling order as before, the electron configuration for Fluorine is: \(1s^{2}2s^{2}2p^{5}\).
03

Determine the hybrid orbitals for B-F bonding in \(\mathrm{BF}_{3}\)

In the \(\mathrm{BF}_{3}\) molecule, the Boron atom forms covalent bonds with three Fluorine atoms. To do this, the Boron atom must have three unpaired electrons in its valence shell. Boron's ground state electron configuration is \(1s^{2}2s^{2}2p^{1}\), which has only one unpaired electron. Thus, the Boron atom must undergo hybridization to obtain 3 unpaired electrons for bonding. During this process, the 2s and 2p orbitals will hybridize, forming three sp2 orbitals. In these sp2 orbitals, the Boron atom will have 3 unpaired electrons, which can form bonds with the 2p electrons of the Fluorine atoms.
04

Identify any unhybridized valence orbitals on B in \(\mathrm{BF}_{3}\)

As we determined in the previous step, the Boron atom forms three sp2 hybrid orbitals for bonding with the Fluorine atoms in the \(\mathrm{BF}_{3}\) molecule. In the hybridization process, one of the 2p orbitals remains unhybridized. Therefore, the 2p_z orbital on Boron is unhybridized in the \(\mathrm{BF}_{3}\) molecule.

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Most popular questions from this chapter

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 \(\pi_{2 \mathrm{p}}\) 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.

Many compounds of the transition-metal elements contain direct bonds between metal atoms. We will assume that the \(z\) -axis is defined as the metal-metal bond axis. (a) Which of the 3 d orbitals (Figure 6.23 ) is most likely to make a \(\sigma\) bond between metal atoms? (b) Sketch the \(\sigma_{3 d}\) bonding and $\sigma_{3 d}^{*}$ antibonding MOs. (c) With reference to the "Closer Look" box on the phases oforbitals, explain why a node is generated in the \(\sigma_{3 d}^{*}\) MO. (d) Sketch the energylevel diagram for the \(\mathrm{Sc}_{2}\) molecule, assuming that only the \(3 d\) orbital from part (a) is important. (e) What is the bond order in \(\mathrm{Sc}_{2} ?\)

Consider the Lewis structure for acetic acid, which is known as vinegar: CCC(=O)O (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 oxygen atoms, and what are the approximate bond angles at the oxygen that is connected to carbon and hydrogen? (c) What is the total number of \(\sigma\) bonds in the entire molecule, and what is the total number of \(\pi\) bonds?

Consider the \(\mathrm{H}_{2}^{+}\) ion. (a) Sketch the molecular orbitals of the ion and draw its energy-level diagram. (b) How many electrons are there in the \(\mathrm{H}_{2}^{+}\) ion? (c) Write the electron configuration of the ion in terms of its MOs. (d) What is the bond order in \(\mathrm{H}_{2}^{+}\) ? (e) Suppose that the ion is excited by light so that an electron moves from a lower-energy to a higher- energy MO. Would you expect the excited-state \(\mathrm{H}_{2}^{+}\) ion to be stable or to fall apart? (f) Which of the following statements about part (e) is correct: (i) The light excites an electron from a bonding orbital to an antibonding orbital, (ii) The light excites an electron from an antibonding orbital to a bonding orbital, or (iii) In the excited state there are more bonding electrons than antibonding electrons?

Benzaldehyde, \(\mathrm{C}_{7} \mathrm{H}_{6} \mathrm{O}\), is a fragrant substance responsible for the aroma of almonds. Its Lewis structure is O=Cc1cccc(C=O)c1 (a) What is the hybridization at each of the carbonatoms of the molecule? (b) What is the total number of valence electrons in benzaldehyde? (c) How many of the valence electrons are used to make \(\sigma\) bonds in the molecule? (d) How many valence electrons are used to make \(\pi\) bonds? (e) How many valence electrons remain in nonbonding pairs in the molecule?
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