Final answer:
In a lead atom, an electron experiences the greatest effective nuclear charge in the 1s orbital, which is the closest to the nucleus. The 1s electrons are less shielded and hence feel the nuclear charge more strongly than electrons in outer orbitals.
Step-by-step explanation:
The electron in a lead atom experiences the greatest effective nuclear charge (Zeff) in the orbital closest to the nucleus. Since energy and orbital radius decrease as nuclear charge increases, the most stable orbital with the lowest energy for an electron is the one nearest to the nucleus. In lead (Pb), which has an electron configuration that ends in 6p², the electrons in the innermost orbitals, such as the 1s orbital, will experience the greatest Zeff. This is because the inner electrons are less shielded by other electrons from the pull of the positively charged nucleus.
The concept of electron shielding indicates that as we move to orbitals farther from the nucleus, the effective nuclear charge felt by these electrons decreases due to the shielding effect of the inner electron shells. Thus, electrons in the outer orbitals, such as the valence electrons, experience less nuclear charge. For the case of lead, as we progress from the 1s to higher shells like 6s or 6p, the Zeff decreases. Therefore, the electron in the 1s orbital of lead would experience the greatest Zeff.