Final answer:
Ionization energy increases across a period because of the rising effective nuclear charge and decreases down a group due to increased electron shielding and the addition of energy levels.
Step-by-step explanation:
Ionization energy, the amount of energy required to remove an electron from an atom, exhibits specific periodic trends:
It generally increases across a period due to a rise in effective nuclear charge. This means that as you move from left to right across the periodic table, electrons are held more tightly by the nucleus because the number of protons increases, leading to a greater attraction.
It decreases down a group because of increased shielding by inner electrons and the addition of valence electron shells, which increase the distance between the nucleus and the valence electrons, thus reducing the nuclear pull on these electrons.
These variations are due to two main factors: the effective nuclear charge that the valence electrons experience and the degree of shielding from the nucleus provided by inner electrons. Therefore, the correct statement is that ionization energy increases across a period due to increasing effective nuclear charge (a), and ionization energy decreases down a group (d), but not due to increasing effective nuclear charge; rather, it decreases due to increased shielding and additional energy levels.