What is Hund’s Rule? Definition, Importance & Examples
Hund’s Rule is essential in chemistry and helps students understand atomic structure, electronic configuration, and magnetic properties of elements.
By knowing Hund’s Rule, you can better grasp how electrons fill up orbitals and why this affects the stability and reactivity of atoms. Let’s break down this important concept in a simple way for all learners.
What is Hund’s Rule in Chemistry?
A Hund’s Rule refers to the guideline for how electrons occupy degenerate (same energy) orbitals within a subshell. This concept appears in chapters related to atomic structure, quantum numbers, and electron configuration, making it a foundational part of your chemistry syllabus.
Hund’s Rule states: Electrons will fill all available degenerate orbitals singly before any begins to pair up. This lowers electron repulsion and increases atom stability.
Molecular Formula and Composition
Hund’s Rule is a rule, not a molecule, so it does not have a molecular formula. It is, however, closely related to the arrangement of electrons in s, p, d, and f orbitals, covering elements from hydrogen to the heaviest metals in the periodic table.
Preparation and Synthesis Methods
There are no chemical synthesis methods for Hund’s Rule, since it is a principle used in assigning electron configurations. Instead, we “apply” Hund’s Rule whenever we write out how electrons fill orbitals in atoms.
Physical Properties of Hund’s Rule
- Being a rule, Hund’s Rule has no physical properties. Nevertheless, it predicts that atoms with unpaired electrons (due to this rule) often show paramagnetic behavior, meaning they are attracted to magnets.
- This contrasts with atoms having all paired electrons, which are diamagnetic and slightly repelled by magnets.
Chemical Properties and Reactions
Hund’s Rule impacts the chemical properties of an atom by affecting its electronic configuration. For example, the stability and magnetic properties of elements depend on how many unpaired electrons remain after following Hund’s Rule.
Frequent Related Errors
- Pairing electrons in orbitals before all degenerate orbitals are singly occupied.
- Ignoring electron spin direction when filling orbitals singly.
- Confusing Hund’s Rule with the Aufbau Principle (order of filling) or Pauli Exclusion Principle (two electrons per orbital, opposite spins).
Uses of Hund’s Rule in Real Life
- Hund’s Rule is widely applied in chemistry and physics while determining the electron configuration of atoms and ions.
- This rule helps explain why elements like oxygen and iron show magnetic properties and why certain elements are chemically reactive.
- It also assists in understanding periodic trends and predicting the magnetic behavior of different materials.
Relation with Other Chemistry Concepts
Hund’s Rule is closely related to topics such as Atomic Orbitals, Electronic Configurations, Aufbau Principle, and Pauli Exclusion Principle. These connections create a strong foundation in atomic theory and prepare students for advanced chemistry and physics concepts.
Step-by-Step Reaction Example
Here’s how Hund’s Rule is applied using the carbon atom (atomic number 6) as an example:
1. Write the total number of electrons: 62. Distribute electrons in orbitals following Aufbau and Pauli Exclusion Principles:
3. Now apply Hund’s Rule to the 2p electrons:
Final Outcome: In the 2p subshell, you will have two unpaired electrons, each in a separate orbital.
Lab or Experimental Tips
Remember Hund’s Rule by the rule of "single first, double later": In any set of boxes (orbitals), always put one arrow (electron) in each box before putting two in one box. Vedantu educators use arrow diagrams with boxes to help students visualize and avoid mistakes during orbital filling.
Try This Yourself
- Draw the orbital filling diagram for nitrogen (atomic number 7) following Hund’s Rule.
- Identify the number of unpaired electrons in oxygen and state if it's magnetic.
- State the difference between Hund’s Rule and Pauli Exclusion Principle in one sentence each.
Final Wrap-Up
We explored Hund’s Rule—its definition, how it controls electron filling in orbitals, and its real-life importance in predicting magnetism and atomic stability. For more stepwise explanations and visual guides check, Aufbau Principle, and other atomic theory resources on Vedantu.
FAQs on Hund’s Rule in Chemistry: Explained with Examples
1. What is Hund’s Rule in Chemistry?
Hund’s Rule states that electrons occupy every degenerate orbital singly before any orbital is doubly occupied, to maximize multiplicity and stability.
Key Points:
- Applies to degenerate orbitals within a subshell (such as p, d, or f orbitals).
- Ensures electrons remain unpaired whenever possible for increased atomic stability.
2. How does Hund’s Rule work for electron configuration?
Hund’s Rule guides the arrangement of electrons in degenerate orbitals by instructing that:
- Each orbital receives one electron before any pairing begins.
- Electrons enter orbitals with parallel spins to minimize repulsion.
- This results in the maximum number of unpaired electrons within a subshell.
3. Can you give an example of Hund’s Rule?
Example: For nitrogen (N, atomic number 7), the electron configuration is 1s2 2s2 2p3.
In the 2p subshell, each of the three p orbitals contains one unpaired electron, all with parallel spins. This follows Hund’s Rule and results in maximum stability.
4. Why is Hund’s Rule important in Chemistry?
Hund’s Rule is crucial because it:
- Ensures maximum unpaired electrons, increasing atom stability.
- Reduces electron-electron repulsion by avoiding premature pairing.
- Helps predict magnetic properties—more unpaired electrons mean higher paramagnetism.
5. What is the difference between Hund’s Rule, Aufbau Principle, and Pauli Exclusion Principle?
Key Differences:
- Aufbau Principle: Electrons fill orbitals in order of increasing energy.
- Hund’s Rule: Electrons singly occupy all degenerate orbitals before pairing.
- Pauli Exclusion Principle: No two electrons in the same atom can have the same set of quantum numbers (must have opposite spins in one orbital).
6. Where is Hund’s Rule applied in electron configuration?
Hund’s Rule is applied when filling degenerate orbitals within a subshell, such as:
- p orbitals (three: px, py, pz)
- d orbitals (five: dxy, dyz, dzx, dx2-y2, dz2)
- f orbitals (seven)
7. Does Hund’s Rule have any exceptions?
Hund’s Rule is generally consistent for neutral atoms, but in rare cases involving complex ions, molecules, or unusual external influences, other effects may override it. For most atom-based electron configurations, Hund’s Rule holds true and is followed closely.
8. How does Hund’s Rule affect magnetic properties?
Hund’s Rule results in unpaired electrons, which lead to paramagnetism in atoms or ions.
- More unpaired electrons mean stronger paramagnetic behavior.
- Fully paired electrons result in diamagnetism (no net magnetism).
9. What happens if Hund’s Rule is ignored in electron configuration?
If Hund’s Rule is not followed:
- Electron configurations will be incorrect.
- Predictions about an element’s chemical and magnetic properties may also be wrong.
- Exam answers and practical applications would be inaccurate.
10. How is Hund’s Rule shown visually in diagrams?
Hund’s Rule is illustrated using orbital box diagrams:
- Each box represents an orbital.
- Arrows (↑ or ↓) show individual electrons.
- Boxes are filled singly (with arrows in the same direction) before pairing starts.
11. Does Hund’s Rule apply to all types of orbitals?
Hund’s Rule specifically applies to degenerate orbitals within the same subshell (e.g., all three p orbitals, all five d orbitals). It does not apply across orbitals of different energy levels or types.
12. What are common mistakes students make with Hund’s Rule?
Common mistakes include:
- Pairing electrons in an orbital before all degenerate orbitals are singly occupied.
- Confusing Hund’s Rule with Aufbau Principle or Pauli Exclusion Principle.
- Incorrectly drawing orbital diagrams, leading to wrong electron configurations and grading errors.
