As you go down the first group (or any group), you are adding protons and electrons. As you add more and more protons, the effective nuclear charge will increase.
atomic radius
Electron shielding decreases the effective nuclear charge.
1.5
inner electrons
2
atomic radius
Electron shielding decreases the effective nuclear charge.
Atomic radius decreases horizontally in periodic table. This is due to increase in nuclear charge.
moving from left to right across a period, one electron is added for each element.example: Boron has 3, Carbon has 4.
Because as the nuclear charge increases, the attraction between the nucleus and the electrons increases and it requires more energy to remove the outermost electron and that means there is a higher ionization energy. As you go across the periodic table, nuclear charge is the most important consideration. So, going across the periodic table, there should be an increase in ionization energy because of the increasing nuclear charge.
the effective nuclear charge on barium is 2.
1.5
Effective nuclear charge is the net charge of an electron in an atom.Z(eff) = Z - S where:Z - atomic numberS - number of shielding electrons
The atomic radius decreases along a period. It is because of increasing effective nuclear charge along a period.
Atomic radius increases down a group due to increase in number of shells. Its value decreases along a period due to increase in nuclear charge.
as you go from left to right across the periodic table, you are adding one more proton every time hence increasing the nuclear charge of the atom. An increased nuclear charge means the nucleus attracts the electrons more to it hence the size of the atom decreases as it "tightens". hence the atomic radius decreases every time.
Atomic radii decreases on moving from left to right as the effective nuclear charge increases.