Give the ground state electron configuration for \( O^{2} \) \( 1 s^{2} 2 s^{2} 2 p^{2} \) \( 1 s^{2} 2 s^{2} 2 p^{5} \) \( 1 s^{2} 2 s^{2} 2 p^{5} \) \( 1 s^{2} 2 s^{2} 2 p^{3} \) \( 1 s^{2} 2 s^{2} 2 p^{4} \)

To determine the ground state electron configuration for the oxide ion \( O^{2-} \), we first need to know the electron configuration of a neutral oxygen atom, which has 8 electrons. The electron configuration for neutral oxygen is: \[ 1s^2 2s^2 2p^4 \] When oxygen gains 2 electrons to form the \( O^{2-} \) ion, these additional electrons will fill the 2p subshell. Therefore, the electron configuration for \( O^{2-} \) becomes: \[ 1s^2 2s^2 2p^6 \] However, this specific configuration is not listed in your options. The closest option that represents the electron configuration of neutral oxygen is: \[ 1s^2 2s^2 2p^4 \] So, the correct answer from the options provided is: \[ 1s^2 2s^2 2p^4 \] This represents the ground state electron configuration of neutral oxygen, but keep in mind that \( O^{2-} \) would actually have \( 1s^2 2s^2 2p^6 \). If you are looking for the configuration of \( O^{2-} \), it would be \( 1s^2 2s^2 2p^6 \), but since that is not an option, the configuration for neutral oxygen is the closest match.