Explain the concept of electron shielding.

Electron shielding refers to the decrease in attraction between an electron and the nucleus due to the presence of other electrons.

In an atom, electrons are arranged in energy levels around the nucleus. The nucleus, being positively charged, attracts the negatively charged electrons. However, the electrons in the inner energy levels, or shells, can shield or block the pull of the nucleus on the outer electrons. This phenomenon is known as electron shielding.

The concept of electron shielding is crucial in understanding the atomic structure and periodic trends. It explains why the atomic radius increases down a group in the periodic table. As we move down a group, the number of energy levels increases. Each new energy level is further from the nucleus than the last and is shielded by the inner energy levels. Therefore, the outermost electrons experience less attraction to the nucleus, causing the atomic radius to increase.

Electron shielding also affects ionisation energy, which is the energy required to remove an electron from an atom. The greater the shielding, the less energy is needed to remove an outer electron, as it is less attracted to the nucleus. This is why ionisation energy decreases down a group.

However, electron shielding is not a perfect block. The shielding effect is less efficient when the shielding electrons and the outermost electron are not in the same type of orbital. This is due to the different shapes and orientations of the orbitals, which can allow the nucleus's pull to reach the outermost electron to some extent.

In summary, electron shielding is a key concept in atomic structure and periodicity. It describes how inner electrons can reduce the effective nuclear charge experienced by outer electrons, influencing atomic size and ionisation energy. Understanding this concept can help you predict and explain trends in the periodic table.