What is shielding effect?
How is it related to periodic table?
How is it related to periodic table?
Briefly - Shielding: inner electrons tend to shield the outer electrons from the attractive force of the nucleus. The more energy levels between the valence electrons and the nucleus, the more shielding.
Lastly, the nuclear charge increases as you go across and down the table, while the shielding stays constant across the periods, but increases as you go down the columns. These tendencies tell you about atom size. Atoms and ions get bigger as you go down the columns because the shielding effect outweighs the effects of the nuclear charge, so the attraction between the nucleus and electrons is weaker and the atom expands in size. In contrast, atoms get smaller as you go across the periods because the nuclear charge effect outweighs the shielding effect, so the attraction between the nucleus and the electron is greater and the atom shrinks in size. Explanation The shielding effect refers to the reduced attraction between outer shell electrons (which are negatively charged) and the positively charged nucleus of an atom. As you descend a group in the periodic table the shielding effect increases as more electrons are added to the outer shells and more shells are filled increasing the distance from the nucleus and reducing nuclear charge. As you go along the periodic table the opposite happens the shielding effect decreases and the nuclear charge to valence electrons increases. Trend in periodic table The reason d-orbitals make a difference is that electrons in d-orbitals do not screen nuclear charge as effectively as those in s and p orbitals. This is because of something called penetration .. essentially the mathematical shapes of d-orbitals prevent them from allowing electrons to penetrate very closely to the nucleus, compared with electrons in s or p-orbitals. In gallium, you have 10 electrons in the filled 3d-subshell, and each of these electrons is doing a slightly worse job (relatively speaking) of screening the nuclear charge than the electrons in the s and p orbitals. Therefore, the effective nuclear charge in gallium is slightly higher than that in aluminum, so the increase in the radius is a quite a bit smaller than would be expected based on the difference between boron and aluminum, or gallium and indium. The trend goes 82 pm (B) -- 118 pm (Al) -- 126 pm (Ga) -- 144 pm (In) [covalent radii from www.webelements.com]. This effect is generally known as the "d-block contraction". (It can be more or less pronounced depending on how you define the atomic radii.)
A similar thing happens (in principle) when you go from indium to thallium ... except in this case you are now dealing with adding a filled f-subshell to the valence shell. Electrons in f-orbitals are even worse at screening nuclear charge than those in d-orbitals, therefore again, the effective nuclear charge in thallium is a bit larger than it is in indium, so again the jump in radius is fairly small (from 144 to 148 pm). This effect (of the filled f-subshell) is generally known as the lanthanide contraction
- Shielding: inner electrons tend to shield the outer electrons from the attractive force of the nucleus. The more energy levels between the valence electrons and the nucleus, the more shielding.
The reason d-orbitals make a difference is that electrons in d-orbitals do not screen nuclear charge as effectively as those in s and p orbitals. This is because of something called penetration .. essentially the mathematical shapes of d-orbitals prevent them from allowing electrons to penetrate very closely to the nucleus, compared with electrons in s or p-orbitals. In gallium, you have 10 electrons in the filled 3d-subshell, and each of these electrons is doing a slightly worse job (relatively speaking) of screening the nuclear charge than the electrons in the s and p orbitals. Therefore, the effective nuclear charge in gallium is slightly higher than that in aluminum, so the increase in the radius is a quite a bit smaller than would be expected based on the difference between boron and aluminum, or gallium and indium. The trend goes 82 pm (B) -- 118 pm (Al) -- 126 pm (Ga) -- 144 pm (In) [covalent radii from www.webelements.com]. This effect is generally known as the "d-block contraction". (It can be more or less pronounced depending on how you define the atomic radii.)
A similar thing happens (in principle) when you go from indium to thallium ... except in this case you are now dealing with adding a filled f-subshell to the valence shell. Electrons in f-orbitals are even worse at screening nuclear charge than those in d-orbitals, therefore again, the effective nuclear charge in thallium is a bit larger than it is in indium, so again the jump in radius is fairly small (from 144 to 148 pm). This effect (of the filled f-subshell) is generally known as the lanthanide contraction
