Question

Why does ionization energy decrease as you move down a group in the periodic table?

1. because shielding decreases
2. because the atoms get closer to noble gas configuration
3. because nuclear charge decreases
4. because the outer electrons get further from the nucleus
5. because effective nuclear charge increases

Answer 1

1.Due to an increase in the shielding effect, and
4.Due to an increase in the atomic radius.

Explanation:

Let's go through the choices one by one:

1. Shielding effect
The shielding effect refers to the repulsion between electron(s) on the valence shell and the inner shell(s).

These repulsions reduce the effect of the electrostatic attraction between the electrons and the positively-charged nucleus, making their removal easier.

2. Valence shell configuration
Notice that all of the elements in the same group (saving for helium of group eighteen, which has only `1s^2`) have an identical valence shell configuration. Meaning that each of them would need to lose (or gain, depending on the group) the same number of electrons to achieve the noble gas electron configuration.

3. Nuclear charge
As the proton number increases, the nuclear charge increases as you move down a period.

Notice that because valence electrons tend to lie so far away from the nucleus, the large separation would outweigh the high nuclear charges and in effect reduces the nucleus' electrostatic grasp on its valence electrons.

However, ionizing energies of the inner shell electrons do tend to increase as you move down a group.

4. Distance between outer shell electrons and the atomic nuclei
This choice can be correct since as the atomic number increases new electrons are added to orbits and orbitals of increasing energy- and therefore increasing distances from the atomic nuclei. This factor, in addition to the shielding effect due to inner shell electrons, reduces the net electrostatic force acting on the valence shell electrons and as a result reduces their ionizing energy.

5. Effective nuclear charge
The effective nuclear charge is "the charge experienced by a specific electron within an atom," (CK-12 Science) it can be approximated with the equation `Z_("eff")=Z-S` where `Z` is the atomic number and `S` is the number of inner shell electrons.

However, this approximation isn't likely to work when comparing elements in the same group since the equation would always end up giving the group number.

Sources:
"The Shielding Effect and Effective Nuclear Charge", Luman Learning, https://courses.lumenlearning.com/introchem/chapter/the-shielding-effect-and-effective-nuclear-charge/
"Electron Shielding", CK-12 Foundation, https://www.ck12.org/chemistry/electron-shielding/lesson/Electron-Shielding-CHEM/
"Periodic Trends in Ionization Energies", CK-12 Foundation, https://www.ck12.org/book/CK-12-Chemistry-Second-Edition/r18/section/9.4/