The concept used in the above question is the trend in first ionization energy across the periodic table. The first ionization energy is the energy required to remove one electron from a neutral atom in its gaseous state. The trend in first ionization energy can be explained by the periodic trends of atomic radius, effective nuclear charge, and electron shielding.
First, as you move from left to right across a period of the periodic table, the atomic radius generally decreases, and the effective nuclear charge (the attraction of the positively charged nucleus for the negatively charged electrons) increases. This makes it more difficult to remove an electron, which means that the first ionization energy generally increases.
Second, as you move down a group of the periodic table, the atomic radius generally increases, and the electron shielding (the repulsion between electrons in different energy levels) increases. This makes it easier to remove an electron, which means that the first ionization energy generally decreases.
Using these trends, we can arrange the given elements in order of increasing first ionization energy, as shown in the previous an
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