Final answer:
Atomic size within a group is affected mainly by electron shielding and effective nuclear charge. As you move down a group, the addition of electron shells increases electron shielding, which offsets the increase in nuclear charge, leading to larger atomic radii. This pattern is observed in the increasing sizes of atoms from lithium to potassium in Group 1.
Step-by-step explanation:
The two variables that affect atomic size within a group are electron shielding and effective nuclear charge. As you move down a group in the periodic table, the number of energy levels (shells) in an atom increases, which leads to an increase in atomic size. This increase in energy levels means that inner shell electrons can shield the outermost electrons from the full charge of the nucleus, a phenomenon known as electron shielding.
Moreover, while effective nuclear charge - the net positive charge experienced by an electron in a multi-electron atom - increases due to more protons, the effect of electron shielding is stronger. Consequently, the increased distance between the outermost electrons and the nucleus due to additional electron shells outweighs the increased nuclear charge, resulting in larger atomic radii.
For example, within Group 1 of the periodic table, a lithium (Li) atom is smaller than a sodium (Na) atom, which in turn is smaller than a potassium (K) atom, and so on down the group. This pattern holds true because there are more electron shells as you move from Li to K, thereby increasing electron shielding, which effectively causes the atomic size to increase.