My vague recollection from college chemistry is that the arrangement of inner electrons shields the positive nucleus and sort of neutralizes the positive charge of the protons. If the electrons can't arrange themselves to block all the positive charge, it leaves positive charges exposed and capable of attracting electrons from other atoms. So for oxygen, the inner electrons can block all but 2 of the positive charges, which makes it capable of attracting 2 electrons from other atoms. I think that's the basic idea, but I don't have a doctorate, so the explanation might not be entirely correct.
@@richardlouiechemistrylectures Thank you very much! It is this, right? en.wikipedia.org/wiki/London_dispersion_force I drew it like this (not 3D, orbitals not correct, initially had 2 more e- but might help visual learnes): sketchtoy.com/69176796
How does radius DECREASE as you go from left to right on the chart? I do understand from bottom to top radius decreasing but not from left to right... have I missed something is this information in a previous lecture? Thanks!
Moving from left to right, the number of protons in the nucleus increases. Thus, there is more positive charge pulling on the negatively charged electrons in the outer energy level. The electrons move closer to the nucleus, and the radius decreases.
I don't know if there is really any logic behind the stability of half-filled p levels. It's probably something we just observe. We notice that elements in the nitrogen group are a little more stable, and we attribute the increase in stability to them all having the same electron configuration (half filled p-levels).
Hey sir, I really appreciate your efforts in explaining everything we need in chemistry , do you help in organic chemistry too?or does anyone know someone who explain it in an easy way here?
Yes, shielding affects ionization energy. The more the outer electrons are shielded from the nucleus, the greater the radius of the atom. A large radius means the outer electrons are not held as tightly, so the ionization needed to remove the outer electrons decreases. Hence, greater shielding lowers ionization energy.
Your a blessing in disguise ✌
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Electron affinity is the reason (aka driving force) why e.g. oxygen wants to get 2 e-? So it's basically the atom's wish to have a full sublevel?
My vague recollection from college chemistry is that the arrangement of inner electrons shields the positive nucleus and sort of neutralizes the positive charge of the protons. If the electrons can't arrange themselves to block all the positive charge, it leaves positive charges exposed and capable of attracting electrons from other atoms. So for oxygen, the inner electrons can block all but 2 of the positive charges, which makes it capable of attracting 2 electrons from other atoms. I think that's the basic idea, but I don't have a doctorate, so the explanation might not be entirely correct.
@@richardlouiechemistrylectures Thank you very much! It is this, right? en.wikipedia.org/wiki/London_dispersion_force
I drew it like this (not 3D, orbitals not correct, initially had 2 more e- but might help visual learnes): sketchtoy.com/69176796
👍👍👍
@@richardlouiechemistrylectures
Thank you sir
Your way of teaching ,wow!
Nice sir you are so genius
Quick Question, so would that mean elements in Group 7 & 15 also have an Elevated IE because they have a half full and full d-sub level?
How does radius DECREASE as you go from left to right on the chart?
I do understand from bottom to top radius decreasing but not from left to right... have I missed something is this information in a previous lecture?
Thanks!
Moving from left to right, the number of protons in the nucleus increases. Thus, there is more positive charge pulling on the negatively charged electrons in the outer energy level. The electrons move closer to the nucleus, and the radius decreases.
I didn't get how electron affinity becomes zero in N, in other words, I didn't get the logic behind half filled P - sub level and electron affinity..
I don't know if there is really any logic behind the stability of half-filled p levels. It's probably something we just observe. We notice that elements in the nitrogen group are a little more stable, and we attribute the increase in stability to them all having the same electron configuration (half filled p-levels).
Hey sir, I really appreciate your efforts in explaining everything we need in chemistry , do you help in organic chemistry too?or does anyone know someone who explain it in an easy way here?
@@lazarl7699 yes
@@noblestarleadership Did you find anyone?
what about shielding, does it affect ionization energy?
Yes, shielding affects ionization energy. The more the outer electrons are shielded from the nucleus, the greater the radius of the atom. A large radius means the outer electrons are not held as tightly, so the ionization needed to remove the outer electrons decreases. Hence, greater shielding lowers ionization energy.
@@richardlouiechemistrylectures thanks
Genious
"'horizontal columns" Ahh yes, an expert educator
Yeah, I should have said "horizontal rows."
Lol, I bet Nicholas is proud of himself. Great job Nicholas, you are bright enough to call out what everyone else understood implicitly. 😂