Final answer:
Electrons in the same valence shell do not effectively shield each other, affecting the effective nuclear charge (Zeff) slightly. As a result, the Zeff increases as we move across a period in the periodic table, leading to a stronger pull on electrons towards the nucleus.
Step-by-step explanation:
Electrons in the same valence shell do not screen one another very effectively, but they do affect the value of effective nuclear charge (Zeff) slightly. While electrons in filled inner shells are highly effective at shielding electrons in outer shells from the nuclear charge, the electrons in the same principal shell are less efficient at this due to the shielding being determined by the likelihood of an electron being between the electron of interest and the nucleus, coupled with electron-electron repulsions.
This has an impact on the properties of elements across the periodic table. As we move from left to right across a period, the nuclear charge (Z) increases by one with each element, but since the electrons in the same valence shell are not effective at shielding, the overall shielding effect increases only a bit. Consequently, the Zeff also increases, which means a stronger pull on the electrons towards the nucleus, making covalent radii smaller and affecting the atom's electronegativity.