What is the ground-state electron configuration of the oxide ion O2−O2−?
Express your answer in condensed form, in order of increasing orbital energy. For example, [He]2s22p2[He]2s22p2 should be entered as [He][He]2s^22p^2 .
The Correct Answer and Explanation is:
The ground-state electron configuration of the oxide ion (O²⁻) can be derived by first understanding the electron configuration of a neutral oxygen atom and then considering the effect of the extra electrons gained when it becomes O²⁻.
Step-by-step Explanation:
- Neutral Oxygen Atom (O): Oxygen has an atomic number of 8, meaning it has 8 protons and, in its neutral state, 8 electrons. The electron configuration of oxygen (in its ground state) is:
[
1s^2 2s^2 2p^4
]
This indicates that oxygen has 2 electrons in the 1s orbital, 2 electrons in the 2s orbital, and 4 electrons in the 2p orbitals. - Oxide Ion (O²⁻): The oxide ion, O²⁻, is formed when the oxygen atom gains 2 additional electrons. These extra electrons will enter the next available orbitals in increasing energy. Since the 1s and 2s orbitals are already full, the two additional electrons will go into the 2p orbitals. After adding 2 electrons, the electron configuration of O²⁻ becomes:
[
1s^2 2s^2 2p^6
]
This is the configuration for a full octet of electrons in the second energy level. - Condensed Electron Configuration: The electron configuration of O²⁻ is now the same as that of neon (Ne), which is a noble gas with a full electron configuration. In condensed form, we write:
[
[Ne] 2s^2 2p^6
]
Since the 2p orbitals are fully occupied with 6 electrons, the oxide ion has a stable, low-energy electron configuration.
Conclusion:
The ground-state electron configuration of the oxide ion (O²⁻) is:
[
[Ne] 2s^2 2p^6
]
This represents a stable configuration where the oxide ion has a full second energy level (2s and 2p orbitals), and it shares the same electron configuration as the noble gas neon (Ne), which is chemically stable due to its full valence shell.