The electronic configuration of nitrogen is `1s^2,2s^(2)2p_x^(1)2p_y^(1)2p_z^(1)` and not `1s^(2),2s^(2)2p_x^(2)2p_(y)^(1)2p_(z)^(0)` . This is explained by

Options

A. Uncertainty principle
B. Aufbau's principle
C. Pauli's exclusion principle
D. Hund's rule

Answer

To understand why the electronic configuration of nitrogen is `1s^2, 2s^2, 2p_x^1, 2p_y^1, 2p_z^1` instead of `1s^2, 2s^2, 2p_x^2, 2p_y^1, 2p_z^0`, we can break down the explanation into several steps based on fundamental principles of quantum mechanics.

### Step-by-Step Solution:

1. **Understanding the Electron Configuration**:
- The correct electronic configuration for nitrogen (atomic number 7) is `1s^2, 2s^2, 2p^3`.
- This means that nitrogen has 2 electrons in the 1s orbital, 2 electrons in the 2s orbital, and 3 electrons in the 2p orbitals.

2. **Pauli Exclusion Principle**:
- According to the Pauli Exclusion Principle, no two electrons in an atom can have the same set of four quantum numbers.
- This implies that each orbital can hold a maximum of two electrons with opposite spins.
- In the case of nitrogen, the three 2p orbitals (2p_x, 2p_y, 2p_z) can each hold one electron before any pairing occurs.

3. **Hund's Rule**:
- Hund's Rule states that electrons will fill degenerate orbitals (orbitals of the same energy) singly and with parallel spins before pairing up.
- For nitrogen, the three 2p electrons occupy the three 2p orbitals (2p_x, 2p_y, 2p_z) singly, which maximizes the total spin and results in a more stable configuration.

4. **Energy Considerations**:
- The Aufbau Principle dictates that electrons fill orbitals starting from the lowest energy level to the highest.
- In nitrogen, the 2p orbitals are of the same energy level, so the electrons will fill each of these orbitals singly before any pairing occurs.

5. **Conclusion**:
- Therefore, the configuration `1s^2, 2s^2, 2p_x^1, 2p_y^1, 2p_z^1` is favored because it follows both Hund's Rule and the Pauli Exclusion Principle, leading to a more stable arrangement of electrons.