To construct the orbital diagram for the sulfide ion (S²⁻), we first need to look at the electronic configuration of sulfur and how it changes when it gains two electrons to form the ion.
1. **Identify the Atomic Number**: The atomic number of sulfur (S) is 16, which means that a neutral sulfur atom has 16 electrons.
2. **Write the Electron Configuration**: The electron configuration for neutral sulfur is:
1s² 2s² 2p⁶ 3s² 3p⁴.
This means sulfur has 2 electrons in the 1s orbital, 2 in the 2s, 6 in the 2p, 2 in the 3s, and 4 in the 3p orbitals.
3. **Account for the Additional Electrons**: The sulfide ion (S²⁻) means that the sulfur atom has gained two additional electrons. This makes the total number of electrons 18.
4. **Determine the New Electron Configuration**: For S²⁻, the electron configuration becomes:
1s² 2s² 2p⁶ 3s² 3p⁶.
This means the 3p orbital has reached its maximum capacity, filling it up to 6 electrons.
5. **Draw the Orbital Diagram**: In the orbital diagram, we represent the orbitals with lines and place arrows for the electrons. Each line represents an orbital, and we use up and down arrows to denote the two electrons in each orbital:
1s: ↑↓ 2s: ↑↓ 2p: ↑↓ ↑↓ ↑↓ 3s: ↑↓ 3p: ↑↓ ↑↓ ↑↓
In summary, the orbital diagram for the sulfide ion S²⁻ illustrates that the additional electrons have filled the 3p orbital completely, leading to a full outer shell, which provides stability to the ion.