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Hi! Does anyone know how to tell the difference between an electron configuration of the ground state and the electron configuration of an excited state? One of the homework problems in section E pertains to this.
When an electron is excited, it can jump from one orbital to another. It will usually jump to a higher energy orbital. For example, if an atom's electron configuration is 1s2 2s2 2p6, its excited configuration could be 1s2 2s1 2p6 3s1. This means that one electron, which was originally in the 2s orbital, jumped from there to the 3s orbital. I hope this helps!
The ground state of an electron is the configuration of electrons such that the atom is at the lowest potential energy. You can determine this by following the examples done in class and following Hund's, Aufbau's, and the Pauli Exclusion Principle to keep adding electrons to the lowest energy orbitals (with a few exceptions). Any other arrangement of electrons is the atom in an excited state, as there is extra energy that causes the electron to move to orbitals with a higher energy.
Another way to figure out excited vs ground state is if you look at the drawn out electron configuration (the one with the arrows ie its given in some of the book problems) and see if if each orbital in a sub shell contains an electron before electrons begin pairing or if a sub shell orbital (ie 3s) has an electron even though the sub shell below it (ie 2p) doesn't have full orbitals. For example, if you looked at the electron configuration of an atom and saw that there was only one electron in the 2s orbital but there was also one electron in the 2p orbital, this atom would be in the excited state since the 2s orbital wasn't full yet.
Like what was mentioned above by others, pay attention to the written electron configuration to notice an electron "jumping" levels. Additionally, when looking at drawn orbital diagrams, making sure that each orbital is filled before moving on to the next ensures that it is in the ground state.
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