/*! This file is auto-generated */ .wp-block-button__link{color:#fff;background-color:#32373c;border-radius:9999px;box-shadow:none;text-decoration:none;padding:calc(.667em + 2px) calc(1.333em + 2px);font-size:1.125em}.wp-block-file__button{background:#32373c;color:#fff;text-decoration:none} Problem 59 How can the paramagnetism of \(\... [FREE SOLUTION] | 91影视

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Chapter 19: Problem 59

How can the paramagnetism of \(\mathrm{O}_{2}\) be explained using the molecular orbital model?

Short Answer

Expert verified
The paramagnetism of O2 can be explained using the molecular orbital model by analyzing its electronic configuration and filling the molecular orbitals. With a total of 12 valence electrons, the electron configuration in molecular orbitals is as follows: 蟽鈧乻(2), 蟽鈧乻*(2), 蟽鈧俿(2), 蟽鈧俿*(2), 蟺2px(1), 蟺2py(1), and the rest are empty. There are two unpaired electrons in the 蟺2px and 蟺2py orbitals, which indicate paramagnetism in O2 according to the molecular orbital model.

Step by step solution

01

Identify the atomic orbitals and number of valence electrons for oxygen

Oxygen has an electron configuration of 1s虏 2s虏 2p鈦�. Since we are combining two oxygen atoms to form O2, we have a total of (2s虏 2p鈦�) 脳 2 = 2s鈦� 2p鈦� valence electrons to fill the molecular orbitals.
02

Determine the molecular orbital diagram for O2

Now, we need to create a molecular orbital (MO) diagram for O2. The diagram should show the atomic orbitals of individual oxygen atoms on the sides and the molecular orbitals in the middle. Make sure you follow the Aufbau principle, Hund's rule, and the Pauli exclusion principle to correctly fill the orbitals. The molecular orbitals include: 1. Two sigma(蟽) orbitals: 蟽鈧乻 and 蟽鈧俿 2. Two sigma(*) anti-bonding orbitals: 蟽鈧乻* and 蟽鈧俿* 3. Two pi(蟺) bonding orbitals: 蟺2px and 蟺2py 4. Two pi(*) anti-bonding orbitals: 蟺2px* and 蟺2py* 5. One sigma(蟽) bonding orbital: 蟽鈧俻z 6. One sigma(*) anti-bonding orbital: 蟽鈧俻z*
03

Fill the molecular orbitals with electrons

Fill the molecular orbitals with the 12 valence electrons of O2, starting from the lowest energy level according to the Aufbau principle. The number of electrons are: 1. 蟽鈧乻: 2 electrons 2. 蟽鈧乻*: 2 electrons 3. 蟽鈧俿: 2 electrons 4. 蟽鈧俿*: 2 electrons 5. 蟺2px and 蟺2py: 4 electrons (2 in each orbital, as per Hund's rule) 6. 蟽鈧俻z: 0 electrons 7. 蟺2px* and 蟺2py*: 0 electrons 8. 蟽鈧俻z*: 0 electrons
04

Analyze the electron configuration for paramagnetism

Observe the electron configuration in molecular orbitals and focus on the 蟺2px and 蟺2py molecular orbitals. As per Hund's rule, there are 2 unpaired electrons in these orbitals (one in each orbital). Since there are unpaired electrons in O2's molecular orbitals, it exhibits paramagnetism. Thus, the molecular orbital model can explain the paramagnetism of O2 due to the presence of unpaired electrons in its 蟺 orbitals.

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