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A neutral oxygen atom has 8 electrons. The atomic structure consists of orbital levels, which are labeled as 1s, 2s, and 2p. The 1s orbital can hold a maximum of 2 electrons, the 2s orbital can hold a maximum of 2 electrons, and the 2p orbital can hold a maximum of 6 electrons. To draw the orbital diagram, follow these three rules: Aufbau's principle, Hund's rule, and Pauli's exclusion principle. 1. Aufbau Principle: Fill the lower energy levels before filling the higher energy levels. 2. Hund's Rule: When filling degenerate orbitals, place one electron into each orbital before pairing the electrons. 3. Pauli's Exclusion Principle: Each orbital can only hold 2 electrons with opposite spins. Following these rules, we will create an orbital diagram for the oxygen atom: \(1s^2 2s^2 2p^4\) Here is the diagram for the neutral oxygen atom: 1s: ↑↓ 2s: ↑↓ 2p: ↑↑↑↓

When an oxygen atom gains two electrons, it transforms into an oxide ion (O^2-). The two additional electrons will fill up the remaining spaces in the 2p orbital. Here is the orbital diagram for the oxide ion (O^2-): 1s: ↑↓ 2s: ↑↓ 2p: ↑↓↑↓↑↓

Adding a third electron to the oxygen atom would require placing the electron into the next available orbital, which is the 3s orbital. However, this process encounters a few obstacles: 1. Electrostatic Repulsion: When trying to add a third electron, the negatively charged electron experiences repulsion from the negatively charged electrons that are already present in the 2p orbitals. This repulsion creates a barrier that prevents the third electron from being added easily. 2. Energy Requirement: As the electron moves away from the nucleus and into higher energy orbitals, it requires more energy to overcome the electrostatic repulsion. Adding a third electron would necessitate a significant amount of energy. Due to the increased electrostatic repulsion and higher energy requirement, it becomes extremely difficult to add a third electron to an oxygen atom.