Respuesta :

wilfredobeng7

Answer:

Shielding effect, Nuclear charge

Explanation:

moving from the left to the right of the periodic table along the same period, electrons enter into the same shells and so shielding effect is fairly constant. The number of protons, however, increase. this increase in protons lead to an increase in nuclear charge and a subsequent increase in effective nuclear charge. As the effective nuclear charge increases, the Valence electrons are more strongly pulled to the nucleus, resulting in a smaller size. anatomic radii, thus, decreases as one moves from the left to the right of the periodic table in groups 3A, 4A, 5A and 6A.

Moving down the group, atomic radii increase because the numbe of protons increase drastically and although the number of shielding electrons increase, it's not enough to hold back the nuclear charge. Effective nuclear charge, thus, increases as one moves down the groups, leading to an increase in atomic size/radius.

pstnonsonjoku

The reason responsible for the variation in atomic radii in group 3A,4A,5A,6A is increasing size of the nuclear charge.

As we move across the period in the periodic table, the size of the nuclear charge is increased because more protons are added but there is no increase in the number of shells.

As such, the outermost electrons experience greater attraction for the nucleus from left to right across the period.

This increased attraction causes the atoms to become smaller across the period.

Hence, across the period, the atomic radii of elements in groups  3A,4A,5A,6A decreases steadily.

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