When transition metal start losing electrons they lose them from the s orbital before the d orbital. Why is this? The iron(II) ion has 24 electrons in this configuration: [Ar] 3d6 The neutral chromium atom also has 24 electrons, but in this configuration: [Ar] 3d5 4s1 I understand that empty, half, and full shells are preferred, but I don't under stand why two atoms with the same number of electrons would have different configurations. I would assume it must be due to the different number of protons, but I don't understand what the reason is. This also would seem to lead to the question of which orbital is higher energy for excited states? If 3d is higher energy than 4s then the chromium atom in an excited state would have its 4s electron jump up to 3d, and then match the iron ion. However, if 4s is higher than 3d I would also expect the iron ion to end up matching the chromium atom. I'm sure this all has perfectly logical explanations, but I evidently did not pick it up in my chemistry class. I'd appreciate it if anyone could explain this, or point me to a link where it is explained well.