Why do atomic sizes decrease across a period?

Atomic sizes decrease across a period due to an increase in nuclear charge without a significant increase in shielding effect.

In more detail, as you move across a period in the periodic table from left to right, the number of protons in the nucleus (the atomic number) increases. This means that the positive charge of the nucleus, also known as the nuclear charge, increases. At the same time, the number of energy levels, or shells, containing electrons does not increase significantly. This is because the additional electrons are being added to the same energy level.

The electrons in the same energy level do not effectively shield each other from the increased nuclear charge. This is known as the shielding effect. The shielding effect is the reduction in effective nuclear charge due to the blocking of attraction between the nucleus and the outer electrons by inner electrons. In a period, the shielding effect remains relatively constant because the electrons are being added to the same energy level.

Therefore, the increased nuclear charge pulls the electrons closer to the nucleus, resulting in a smaller atomic radius. This is why atomic sizes decrease across a period. It's a bit like a tug of war between the nucleus and the electrons. As the nucleus gets more positively charged, it pulls the electrons in closer, making the atom smaller.

Remember, this trend is not always perfect due to the complexities of atomic structure, but it is a good general rule to understand and apply when studying the periodic table.