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Chapter 7: Problem 41

Does the minimization of electron–electron repulsions correlate with Hund’s rule?

Short Answer

Expert verified
Yes, the minimization of electron-electron repulsions correlates with Hund's rule. Hund's rule ensures that electrons occupy degenerate orbitals with parallel spins, minimizing repulsive forces between electrons and leading to a lower energy state for the atom. This ultimately enhances the stability of the atom.

Step by step solution

01

Understanding electron-electron repulsion minimization

Electron-electron repulsion occurs when two electrons come close to each other, which causes repulsive forces according to Coulomb's law. Minimizing electron-electron repulsion is essential to achieve a lower energy state in an atom and, ultimately, its stability.
02

Hund's rule and minimizing repulsion

Hund's rule states that electrons in degenerate orbitals (orbitals with the same energy level, for example, the three p orbitals in a subshell) will first fill these orbitals with parallel spins, one electron per orbital, before starting to pair up with opposite spins. Following Hund's rule allows for minimizing electron-electron repulsion.
03

Parallel spins and minimized repulsion

When electrons occupy different orbitals with parallel spins, their electron clouds are less likely to overlap, minimizing their repulsive forces. This results in a lower energy state for the atom, which is favorable in terms of stability.
04

Hund's rule correlation to minimization of electron-electron repulsions

Yes, the minimization of electron-electron repulsions correlates with Hund's rule. By occupying degenerate orbitals with one electron per orbital and parallel spins, Hund's rule helps to minimize the repulsive forces among electrons in an atom. This leads to lower atomic energy states and increased stability.

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Most popular questions from this chapter

How many unpaired electrons are present in each of the following in the ground state: $\mathrm{O}, \mathrm{O}^{+}, \mathrm{O}^{-}, \mathrm{Os}, \mathrm{Zr}, \mathrm{S}, \mathrm{F}, \mathrm{Ar}$ ?

One of the emission spectral lines for \(\mathrm{Be}^{3+}\) has a wavelength of 253.4 \(\mathrm{nm}\) for an electronic transition that begins in the state with \(n=5 .\) What is the principal quantum number of the lower-energy state corresponding to this emission? (Hint: The Bohr model can be applied to one- electron ions. Don't forget the \(Z\) factor: \(Z=\) nuclear charge \(=\) atomic number.)

Which of the following sets of quantum numbers are not allowed in the hydrogen atom? For the sets of quantum numbers that are incorrect, state what is wrong in each set. a. \(n=3, \ell=2, m_{\ell}=2\) b. \(n=4, \ell=3, m_{\ell}=4\) c. \(n=0, \ell=0, m_{\ell}=0\) d. \(n=2, \ell=-1, m_{\ell}=1\)

The electron affinity for sulfur is more exothermic than that for oxygen. How do you account for this?

Answer the following questions assuming that \(m_{s}\) could have three values rather than two and that the rules for \(n, \ell,\) and \(m_{\ell}\) are the normal ones. a. How many electrons would an orbital be able to hold? b. How many elements would the first and second periods in the periodic table contain? c. How many elements would be contained in the first transition metal series? d. How many electrons would the set of 4\(f\) orbitals be able to hold?
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