What is the group (top to bottom) trend in the first ionization energies?Why? It increases because the nuclear charge increases It decreases because the distance from the nucleus to the valence electrons increases It decreases because the nuclear charge decreases It increases because the distance from the nucleus to the valence electrons decreases

## Step 1: Understand the taskHere, we want to know about the first ionization energies trend as we move from top to bottom in a group on the periodic table.## Step 2: Analyze the trendAs we move down a group on a periodic table from top to bottom, the atomic number and hence the nuclear charge increase because of an additional proton in the nucleus. But, this doesn’t imply an increased first ionization energy. Reasons include an increase in atomic size (or the distance of the outermost electron from the nucleus), much more critical to our analysis, and the weakening effect of electron shielding.## Step 3: Reflect on atomic size and electron shieldingAtomic size tends to increase down a group due to the addition of energy levels. With each step down a group, the outermost electrons are located further from the nucleus—increasingly so due to electron shielding. Greater atomic size, more distant valency shells, and stronger electron shielding lead to a decreased nuclear hold on valence electrons, therefore less energy (lower first ionization energy) is needed to move those electrons.### The formula representing this situation is: That means that as the size of the atom (or the distance from the nucleus to the valence electrons) becomes bigger, the energy required to remove an electron from an atom (first ionization energy) decreased which fits our observed trend.