Understanding an atom’s electron configuration is key to predicting its chemical behavior. Hund’s rule is a fundamental principle that governs this arrangement, stating that electrons must occupy orbitals of equal energy one at a time before pairing up. Identifying which electron configuration violates this rule helps us distinguish an atom’s stable ground state from a less stable, excited state, providing deep insight into its reactivity and magnetic properties.
What is Hund’s Rule of Maximum Multiplicity?
Hund’s rule is a core principle in chemistry that helps determine the most stable arrangement of electrons in an atom. Think of it like passengers getting on a bus; people will take an empty row of seats for themselves before they start sitting next to someone else. Electrons do the same thing in orbitals of equal energy, which are called degenerate orbitals.
The rule specifically states that electrons will fill each degenerate orbital with one electron before any orbital gets a second electron. Furthermore, all the single electrons in these separate orbitals will have the same spin. This arrangement minimizes the natural repulsion between negatively charged electrons, leading to a more stable, lower-energy atom.
The primary goal of this electron-filling strategy is to achieve maximum stability. By keeping electrons as far apart as possible in separate orbitals, the atom exists in its most favorable energetic state, known as the ground state.
How Hund’s Rule Works with Other Principles
Hund’s rule does not operate in isolation. It works together with two other important rules to give a complete picture of an atom’s electron configuration.
The first is the Aufbau principle, which dictates that electrons fill the lowest energy orbitals first. This is why the 1s orbital is filled before the 2s, and the 2s before the 2p. Hund’s rule applies when we get to a subshell with multiple orbitals of the same energy, like the three orbitals in the 2p subshell.
The second is the Pauli exclusion principle. This principle ensures that no two electrons in the same atom can have the exact same set of four quantum numbers. In practical terms, this means that if two electrons do occupy the same orbital, they must have opposite spins (one “spin up” and one “spin down”). Hund’s rule dictates the order of filling, and the Pauli exclusion principle dictates the spins once pairing begins.
How to Spot an Electron Configuration that Violates Hund’s Rule
Identifying a violation of Hund’s rule is straightforward if you know what to look for. A violation occurs when an electron pair is formed in a degenerate orbital before all other orbitals of the same energy are occupied by at least one electron. This configuration represents an excited state, not the ground state, because it has higher energy due to increased electron repulsion.
To check for a violation, follow these simple steps:
- Focus on the electrons in the highest-energy subshell (e.g., the p, d, or f subshell).
- Count the number of orbitals in that subshell (p has 3, d has 5, f has 7).
- Check if any orbital contains a pair of electrons while another orbital in the same subshell is completely empty. If so, it’s a clear violation of Hund’s rule.
An atom in such a state is less stable and will quickly release energy to return to its proper ground state configuration.
Common Examples of Correct vs. Incorrect Configurations
Seeing a direct comparison can make the concept much clearer. Let’s look at the p-block elements Carbon and Nitrogen, which have electrons in the 2p subshell. The 2p subshell has three degenerate orbitals (often labeled 2px, 2py, and 2pz).
The table below shows the correct ground state configuration according to Hund’s rule versus an incorrect, higher-energy configuration that violates it. The arrows represent electrons and their spin direction.
| Element (Electrons in 2p) | Correct Ground State (Follows Hund’s Rule) | Incorrect State (Violates Hund’s Rule) |
|---|---|---|
| Carbon (2p²) | ↑ in 2px, ↑ in 2py | ↑↓ in 2px, empty 2py |
| Nitrogen (2p³) | ↑ in 2px, ↑ in 2py, ↑ in 2pz | ↑↓ in 2px, ↑ in 2py, empty 2pz |
| Oxygen (2p⁴) | ↑↓ in 2px, ↑ in 2py, ↑ in 2pz | ↑↓ in 2px, ↑↓ in 2py, empty 2pz |
In each incorrect example, electrons are paired up unnecessarily, leaving an orbital empty. This creates more electron-electron repulsion than the correct ground state configuration, making the atom less stable.
Why Does Violating Hund’s Rule Matter?
The way electrons are arranged directly impacts an atom’s properties. A configuration that violates Hund’s rule is not just a theoretical mistake; it describes an atom in a physically different state with tangible consequences.
Firstly, it affects atomic stability. An atom that violates Hund’s rule is in a higher-energy, or “excited,” state. This state is inherently unstable, and the atom will tend to release energy (often as light) to fall back to its more stable ground state. This principle is the basis for phenomena like atomic emission spectra.
Secondly, it alters chemical and magnetic properties. The number of unpaired electrons determines whether an atom is paramagnetic (attracted to a magnetic field) or diamagnetic (weakly repelled by it). Following Hund’s rule maximizes the number of unpaired electrons, often making atoms paramagnetic. A violation that reduces the number of unpaired electrons would change this magnetic behavior.
Are There Exceptions to these Rules?
While Hund’s rule itself is consistently applied within a subshell, some elements show what are called “anomalous electron configurations.” These are not truly violations of Hund’s rule but are exceptions to the general Aufbau filling order.
The most famous examples are Chromium (Cr) and Copper (Cu). Based on the Aufbau principle, you would expect Chromium’s configuration to end in 4s²3d⁴. However, its actual ground state configuration is 4s¹3d⁵. This happens because a half-filled d-subshell (d⁵) is exceptionally stable. Similarly, Copper adopts a 4s¹3d¹⁰ configuration instead of 4s²3d⁹ to gain the stability of a completely filled d-subshell.
These anomalies prioritize the overall lowest energy state for the entire atom, showing that the stability gained from having a half-filled or full d-subshell outweighs the energy cost of promoting an electron from the 4s orbital.
Frequently Asked Questions
What is Hund’s rule in simple terms?
Hund’s rule states that when filling orbitals of equal energy, electrons will spread out to occupy each orbital individually before any of them start pairing up. This minimizes repulsion and makes the atom more stable.
How do you know if an electron configuration violates Hund’s rule?
Look at the outermost subshell. If you find an orbital with two electrons (a pair) while another orbital of the exact same energy level is empty, the configuration violates Hund’s rule and is not a ground state.
Can you give an example of a Hund’s rule violation?
For a nitrogen atom (1s²2s²2p³), the correct configuration has one electron in each of the three 2p orbitals. A violation would be placing two electrons in the first 2p orbital, one in the second, and leaving the third 2p orbital empty.
Does violating Hund’s rule make an atom unstable?
Yes, an electron configuration that violates Hund’s rule describes an atom in a higher-energy excited state. This state is less stable than the ground state, and the atom will seek to return to the more stable arrangement.
Why do electrons occupy orbitals singly before pairing?
Electrons are all negatively charged and naturally repel each other. By occupying separate orbitals within the same subshell, they can be farther apart, which minimizes this electrostatic repulsion and results in a lower overall energy for the atom.
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