Final answer:
The concept of electron shielding, stemming from electron-electron repulsions, confirms that electrons in lower-energy orbitals shield electrons in higher-energy orbitals from the full nuclear charge, making the initial statement true.
Step-by-step explanation:
The statement that an electron is repulsed by electrons in lower-energy orbitals blocking it from the full effects of nuclear charge is true. This concept is known as shielding, whereby electrons in lower-energy orbitals or inner shells can shield an electron in a higher-energy orbital from the full positive charge of the nucleus. As a result, the electron feels an effective nuclear charge (Zeff) which is less than the actual charge of the nucleus. Higher-energy orbits are located further from the atomic nucleus, and electrons are naturally repelled from each other due to their negative charge. Moving an electron away from the nucleus requires energy because electrons are attracted to the positive protons in the nucleus.
Not all electrons shield equally. Electrons in the same principal shell are not as effective at shielding each other as electrons in filled inner shells are at shielding electron in outer shells. This is why electrons in orbitals that experience more shielding are less stabilized and thus higher in energy, contributing to the complexity of calculating the potential energy in multielectron atoms or ions.