The ionization energy refers to the amount of energy required to remove an electron from an atom or ion in its gaseous state. When moving across a period in the periodic table from left to right, the ionization energy generally increases. There are a few key factors that contribute to this trend:
Increased nuclear charge: As you move across a period, the number of protons in the nucleus increases. Since electrons are negatively charged, they are attracted to the positively charged nucleus. With a greater nuclear charge, the attractive force on the electrons increases, making it more difficult to remove an electron. Therefore, more energy is required to overcome this increased attraction.
Decreased atomic radius: Moving from left to right across a period, the atomic radius generally decreases. The atomic radius is the distance between the nucleus and the outermost electrons. When the atomic radius is smaller, the outermost electrons are closer to the nucleus, experiencing a stronger attraction. This stronger attraction requires more energy to remove an electron, resulting in an increased ionization energy.
Increased electron shielding: Electron shielding refers to the effect of inner electrons shielding the outer electrons from the full attraction of the nucleus. As you move across a period, the number of inner electrons generally remains constant while the nuclear charge increases. The increased nuclear charge has a stronger pull on the outer electrons, but the inner electrons provide some degree of shielding. However, the shielding effect remains relatively constant across a period, so it does not counteract the increasing nuclear charge significantly. Thus, the ionization energy generally increases.
These combined factors of increased nuclear charge, decreased atomic radius, and relatively constant electron shielding contribute to the trend of increasing ionization energy across a period in the periodic table.