shielding effect
Imagine standing at the back of a crowd at a concert, straining to feel the heat of stage lights. The rows of people in front of you absorb much of the glare, so you feel only a faint warmth. An outer electron in an atom is in the same boat: the inner electrons crowd between it and the nucleus, screening off much of the nuclear pull so the outer electron feels a weakened attraction.
The shielding effect (also called screening) is the reduction in the nucleus's pull on an outer electron caused by the repulsion of the electrons lying between it and the nucleus. Inner electrons partly cancel out the positive nuclear charge, so the outer electron experiences not the full charge but a smaller effective nuclear charge. The more inner electrons there are, the stronger the shielding.
Shielding is the reason outer electrons are held loosely, why atoms grow larger down a group, and why ionisation energy drops as shells pile up. The caveat is that shielding is not perfect or all-or-nothing: electrons in the same shell shield each other only weakly, and s electrons, which dip closer to the nucleus, shield more effectively than p or d electrons in the same shell.
Lithium's single outer electron sits behind two inner ones, which screen most of the +3 nuclear charge, so it feels a net pull of only about +1.3. That heavy shielding is why lithium's outer electron is so easy to remove and why lithium is a soft, reactive metal.
Inner electrons screen the nucleus, weakening its grip on the outer ones.
Shielding and effective nuclear charge are two sides of one coin: effective nuclear charge equals the true charge minus the shielding. Down a group shielding grows; across a period it barely changes while protons keep being added.