Iodine has an atomic number of 53, meaning it has 53 electrons distributed across its electron shells. The electron configuration of iodine is 2, 8, 18, 18, 7. To understand why the fourth shell does not reach its maximum of 18 electrons and instead holds 7 before starting a new shell, we need to delve into the structure of electron shells and how they are filled.
Electron shells are filled based on the principles of quantum mechanics and the energy levels of electrons. The first shell can hold a maximum of 2 electrons, the second can hold a maximum of 8 electrons, and the third shell can hold a maximum of 18 electrons. The fourth shell also can theoretically hold more than 18 electrons, but in practice, it often doesn’t reach this number due to energy level differences and stability requirements.
In iodine’s case, after filling the first three shells with a total of 28 electrons (2 + 8 + 18), there are 25 electrons remaining, which fill the fourth shell. However, the fourth shell is often less stable than the inner shells when it is not fully maximized. Consequently, iodine fills its fourth shell with 7 electrons, rather than the maximum of 18, before starting to fill the fifth shell.
This phenomenon occurs because as electrons are added to higher energy levels, they experience increased repulsion due to the electron-electron interactions and the stability provided by the filled inner shells. Essentially, iodine’s electron configuration reflects a balancing act between shielding effects and the tendency of electrons to occupy lower energy levels first for stability.
In summary, iodine only accommodates 7 electrons in its fourth shell due to stability factors and the energies involved in electron positioning, leading to the initiation of a new shell rather than maximizing the fourth shell’s capacity.