As we go across the periodic table, electron shielding (increases/decreases). As we go down, electron shielding (increases/decreases). Help!?

Electron "shielding".....

What you are talking about is the effective nuclear charge. Consider the outermost electron in an atom. It is both attracted to the positively charged nucleus and repelled by the other electrons between it and the nucleus. It is this electrostatic repulsion which is called "shielding." In the equation below the "shielding" is represented by "S".

Zeff = Z - S...... Zeff is the effective nuclear charge, Z is the actual charge on the nucleus and S is the amount of "shielding". You may recognize Z as the atomic number. The atomic number, Z, is both equal to the number of protons in the nucleus and the relative charge on the nucleus.

As Z increases across a period (horizontal row) on the periodic table the effective nuclear charge increases as "S", the shielding, decreases. Going a group (family; vertical column) the effective nuclear charge increases as the number of protons increases and the shielding decreases.

Read more about the effective nuclear charge here: https://en.wikipedia.org/wiki/Effective_nuclear_ch...