The periodic table is organized based on the atomic number of elements, which corresponds to the number of protons in the nucleus of an atom. In a neutral atom, the atomic number also represents the number of electrons in the atom. Elements are arranged in order of increasing atomic number, and this arrangement provides a systematic way to group elements with similar properties together.
However, there are cases where elements don't strictly follow the order of increasing atomic number. This can be observed in certain instances where the electron configuration of an element results in a more stable arrangement by placing electrons in a different orbital. There are two main reasons for this phenomenon: the filling of electron subshells and the interaction between electrons.
- Filling of electron subshells: Electron subshells are the different energy levels or regions where electrons can be found within an atom. The subshells are labeled as s, p, d, and f. According to the Aufbau principle, electrons fill these subshells in a specific order, starting with the lowest energy level. However, in some cases, an electron may enter a higher energy subshell, resulting in an exception to the strict order of atomic numbers.
For example, one such case is chromium (Cr) and copper (Cu). Chromium has an atomic number of 24, suggesting that its electron configuration should be 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁴. However, in reality, the electron configuration of chromium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d⁵. This deviation occurs because it is energetically favorable for chromium to have a half-filled 3d subshell (3d⁵) and a completely filled 4s subshell. Similarly, copper has an atomic number of 29 but its electron configuration is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ 3d¹⁰, as it is more stable to have a completely filled 3d subshell and a partially filled 4s subshell.
- Electron-electron repulsion: Another factor that can influence the electron configuration and lead to deviations from the atomic number order is the repulsion between electrons. Electrons repel each other due to their negative charges, and this repulsion affects the energy levels and stability of the electron arrangement. In some cases, rearranging electrons can minimize the repulsion and result in a more stable configuration.
For example, in the case of molybdenum (Mo), its atomic number is 42, suggesting an electron configuration of 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s² 4d⁴. However, the actual electron configuration of molybdenum is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s¹ 4d⁵. By placing one electron in the 5s subshell instead of the 4d subshell, the electron-electron repulsion is minimized, resulting in a more stable configuration.
These exceptions in the filling order of electron subshells and the influence of electron-electron repulsion can lead to variations from the strict atomic number order in the periodic table. It is important to note that these exceptions occur in specific cases and do not significantly impact the overall organization and usefulness of the periodic table as a tool for understanding and predicting the properties of elements.