Final answer:
The effective nuclear charge experienced by an electron decreases as it moves further away from the nucleus due to electron shielding, where inner electrons weaken the nucleus's hold on outer electrons. Coulomb's Law explains that the attraction is inversely related to the square of the distance between charges.
Step-by-step explanation:
As electrons move further away from the nucleus, the effective nuclear charge they experience does indeed decrease. This is primarily due to a phenomenon known as electron shielding. Electron shielding occurs when inner electrons effectively block the positive charge of the nucleus, reducing the force of attraction felt by electrons located in outer shells. Though the total nuclear charge (the number of protons) remains constant, the presence of intervening electron shells leads to a diminished effective nuclear charge experienced by electrons in higher energy levels.
Coulomb's Law helps us understand that the force of attraction between two charged particles is inversely proportional to the square of the distance between them. Therefore, as the distance between the nucleus and an electron increases, the attractive force decreases. Atomic radii tend to increase across the periodic table as more electron shells are added, increasing the distance between the outermost electrons and the nucleus.
Moreover, the emission of a particle with a -1 charge from the nucleus during certain processes indicates a change in the atom's charge dynamics, further demonstrating the interplay of charges within an atom.