Electrons never pair, if there are empty orbitals in a given sub-shell. This is

Options

A. Aufbau principle
B. Pauli's exclusion principle
C. Hund's rule of maximum multiplicity
D. Heisenberg's uncertainity principle

Answer

To solve the question, we need to identify which principle explains why electrons do not pair up in orbitals if there are empty orbitals available in a given sub-shell. Here’s a step-by-step breakdown:

### Step 1: Understand the Question
The question states that "Electrons never pair if there are empty orbitals in a given sub-shell." We need to determine which principle or rule this statement aligns with.

### Step 2: Review the Relevant Principles
1. **Aufbau Principle**: This principle describes the order in which electrons fill orbitals. It states that electrons occupy the lowest energy orbitals first.
2. **Pauli Exclusion Principle**: This principle states that no two electrons in an atom can have the same set of four quantum numbers. This means that an orbital can hold a maximum of two electrons, and they must have opposite spins.
3. **Hund's Rule**: This rule states that for degenerate orbitals (orbitals of the same energy), electrons will fill each orbital singly before pairing up. This is the key rule that explains the behavior described in the question.
4. **Heisenberg Uncertainty Principle**: This principle relates to the impossibility of simultaneously knowing the exact position and momentum of an electron, but it does not directly address electron pairing.

### Step 3: Identify the Correct Principle
Among the principles reviewed, **Hund's Rule** is the one that specifically states that electrons will occupy empty orbitals singly before pairing up. Therefore, the statement in the question aligns with Hund's Rule.

### Step 4: Conclusion
The correct answer to the question is that the statement "Electrons never pair if there are empty orbitals in a given sub-shell" is explained by **Hund's Rule**.

### Final Answer
The correct option for this statement would be **Hund's Rule**.

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