Ionization energy is the energy required to remove an electron from a gaseous atom or ion. In general, within a group in the periodic table, ionization energy decreases as you move down the group. This is due to the fact that as you move down a group, the number of electron shells increases, causing the outer electrons to be further from the nucleus and hence less tightly bound.
In this context, let's consider the elements beryllium (Be), calcium (Ca), and strontium (Sr). These elements are all in group 2 of the periodic table, but they are positioned at different periods:
- Beryllium (Be) is in the second period.
- Calcium (Ca) is in the fourth period.
- Strontium (Sr) is in the fifth period.
Following the general trend of ionization energy in a group:
1. Beryllium (Be): Being in the second period and higher up in the group, it has the highest ionization energy among these three elements because its electrons are closest to the nucleus.
2. Calcium (Ca): This is in the fourth period, and hence it has a lower ionization energy than beryllium but higher than strontium.
3. Strontium (Sr): Located in the fifth period, it has the lowest ionization energy because its electrons are farthest from the nucleus compared to the others.
Thus, the correct order of ionization energies from lowest to highest is:
[tex]\[ \text{Sr} < \text{Ca} < \text{Be} \][/tex]
Therefore, the correct relationship of these elements regarding their ionization energy is:
[tex]\[ \text{D. Sr < Ca < Be} \][/tex]
