How does the shielding effect affect trends?

The shielding effect is the electrons between the nucleus and the valence electrons acting as a "shield" - repelling the outer electrons because they have the same charge, lowering the effective nuclear charge.

You can calculate the effective nuclear charge (ENC) by

#Z_(eff) = Z - S#

where #Z_(eff)# is the ENC, #Z# is the atomic number (no. protons) and #S# is the shielding electrons.

The more shielding electrons you have, the lower the ENC, so the less force there is holding onto the outer shell electrons.

If there is less force holding onto valence electrons, then they will be lost more easily, and likewise not gained as easily. Therefore when you move down the left-hand-side of the periodic table, atoms become more reactive - more liable to lost electrons. As you move down the right, however, atoms cannot gain electrons as easily, so as you move down the column of halogens, reactivity decreases.

This is because the left of the table lose electrons to form positive ions, while the right gains electrons to form negative ions.

Also, the lower the ENC due to greater shielding effect, the less electronegative an element will be, so whatever electrons it does have will not be pulled towards it so easily.