Which Orbital-Filling Diagram Violates Hund’s Rule?
Hund’s rule is a fundamental principle in atomic structure that governs how electrons occupy orbitals within a subshell. Named after the German physicist Friedrich Hund, this rule ensures that electrons fill degenerate orbitals (orbitals with the same energy level) in a way that maximizes the number of unpaired electrons. Violating Hund’s rule leads to incorrect electron configurations, which can misrepresent an atom’s chemical behavior and magnetic properties. Understanding which orbital-filling diagram violates Hund’s rule is crucial for students and professionals in chemistry to accurately predict molecular bonding and reactivity.
Understanding Hund’s Rule
Hund’s rule states that:
**Electrons will fill degenerate orbitals singly as much as possible before pairing up., the three p-orbitals or five d-orbitals), each orbital receives one electron with parallel spins before any orbital begins to hold a second electron. **
What this tells us is when filling orbitals in the same subshell (e.Consider this: g. This arrangement minimizes electron-electron repulsion and stabilizes the atom.
The rule is closely tied to the Pauli exclusion principle, which prohibits two electrons in the same orbital from having identical quantum numbers. While the Pauli principle dictates that paired electrons must have opposite spins, Hund’s rule determines the order in which orbitals are filled. Together, these principles form the foundation of the Aufbau principle, which describes the sequence of electron occupation in atoms.
Common Violations of Hund’s Rule
An orbital-filling diagram violates Hund’s rule if it shows electrons paired in a single orbital before all degenerate orbitals are singly occupied. Here are examples of such violations:
Example 1: Improper Pairing in the p Subshell
Consider a carbon atom (atomic number 6) with the electron configuration [He] 2s² 2p². The 2p subshell has three degenerate orbitals. A correct diagram would show two unpaired electrons in separate p-orbitals (e.g., one in 2pₓ and one in 2pᵧ). A violation occurs if both electrons are placed in the same orbital (e.g., 2pₓ), leaving the other orbitals empty. This configuration contradicts Hund’s rule and is less stable than the correct arrangement.
Example 2: Premature Pairing in the d Subshell
For a titanium atom (atomic number 22), the 3d subshell contains five orbitals. In the ground state, titanium’s electron configuration is [Ar] 3d² 4s². A correct diagram shows two unpaired electrons in separate d-orbitals. A violation would involve placing both electrons in the same orbital, resulting in paired spins and fewer unpaired electrons than allowed by Hund’s rule.
Example 3: Ignoring Orbital Energy Levels
In some cases, diagrams may incorrectly prioritize filling higher-energy orbitals over lower-energy ones. To give you an idea, in oxygen (atomic number 8), the 2p subshell has four electrons. A violation would occur if three electrons occupy one orbital (with one paired electron) instead of distributing them as two unpaired electrons in two orbitals and two paired electrons in the third.
How to Identify Violations in Orbital-Filling Diagrams
To spot a violation of Hund’s rule, follow these steps:
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- Still, 4. Now, Verify Unpaired Electrons: make sure each orbital in the subshell has at least one electron before any orbital holds a second. Check the Subshell: Identify the subshell in question (e.Think about it: 3. Because of that, Count the Electrons: Determine how many electrons are present in the subshell. g.Worth adding: , s, p, d, f) and count the number of degenerate orbitals it contains. Spin Alignment: Confirm that all unpaired electrons have parallel spins (arrows pointing in the same direction).
A diagram that fails any of these checks likely violates Hund’s rule. To give you an idea, if a p-subshell with three electrons shows two electrons paired in one orbital and the third in another, this is incorrect. The correct configuration would have one electron in each of two orbitals and one paired electron in the third, with all unpaired electrons aligned in the same direction.
FAQ
Q: Why is Hund’s rule important in chemistry?
A: Hund’s rule explains the stability of certain electron configurations, which influence an atom’s magnetic properties, bonding behavior, and reactivity. As an example, elements with unpaired electrons exhibit paramagnetism, while those with all paired electrons are diamagnetic.
Q: Can Hund’s rule ever be broken?
A: In extreme conditions, such as high-pressure environments or in excited states, temporary violations may occur. On the flip side, the ground-state electron configuration of an atom always follows Hund’s rule to minimize energy.
Q: How does Hund’s rule relate to the Aufbau principle?
A: The Aufbau principle determines the order in which orbitals are filled (1s, 2s, 2p, etc.), while Hund’s rule governs the filling of electrons within a subshell. Both principles work together to establish the correct electron configuration Nothing fancy..
Q: What happens if electrons are paired too early in a subshell?
A: Pairing electrons prematurely increases electron-electron repulsion, raising the atom’s energy and destabilizing it. This violates the principle of minimum energy, which drives electron configurations.
Conclusion
An orbital-filling diagram violates Hund’s rule when it pairs electrons in a single orbital before all degenerate orbitals are singly occupied. Understanding Hund’s rule is not just an academic exercise—it is essential for predicting chemical properties and behaviors in real-world applications, from material science to pharmaceuticals. In real terms, this mistake disrupts the natural tendency of electrons to minimize repulsion and maximize stability. By applying the guidelines outlined above, anyone can identify such violations and ensure accurate electron configurations. Mastering this concept paves the way for deeper insights into atomic structure and molecular interactions.
