The atomic radius generally decreases as electrons are added to the same electron shell across a period because the nuclear charge (number of protons) simultaneously increases. This increased nuclear charge exerts a stronger pull on all the electrons, drawing them closer to the nucleus and thus reducing the overall atomic size.
Here's a more detailed breakdown:
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Increasing Nuclear Charge: As you move from left to right across a period in the periodic table, each element has one more proton in its nucleus than the element before it. This means the positive charge of the nucleus is increasing.
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Constant Shielding (Relatively): Since electrons are being added to the same electron shell, the shielding effect (where inner electrons reduce the effective nuclear charge felt by outer electrons) remains relatively constant. The added electrons do shield each other to some extent, but this shielding effect is less significant than the increasing nuclear charge.
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Greater Attraction: The stronger positive charge of the nucleus exerts a greater attractive force on all the electrons in the atom. This increased attraction pulls the electrons closer to the nucleus.
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Decreased Atomic Radius: Because the electrons are being pulled closer to the nucleus, the overall size of the electron cloud, and thus the atomic radius, decreases.
In Summary:
The dominant factor causing the decrease in atomic radius across a period is the increasing nuclear charge, which attracts the electrons more strongly and pulls them closer to the nucleus. Although the number of electrons is also increasing, their shielding effect is not enough to counteract the increased nuclear attraction.
