#Atomic size trend on periodic table full

#Atomic size trend on periodic table full

So when the two nitrogen atoms, who have 5 electrons in their outer shell, get a full shell of 8 electrons by sharing three electrons with the other, they have a much lower total energy than when they are on their own. The reason for these covalent bonds is that atoms lower their energy when they have full electron shells, and everything wants to be in the lowest energy state possible. Nitrogen gas is N2 - with a triple bond - even though two nitrogen atoms necessarily must have the same electronegativity. The general trend is that electronegativity increases as you go up and to the right on the periodic table. Oxygen and Nitrogen are also quite electronegative. If the atom has a smaller atomic radius, it means the electrons are being pulled close to its nucleus much more (generally)īecause of this, Flourine for example would be an electron 'hog'- it is the most electronegative element on the periodic table and will gladly take away and hold onto an electron. For example, Caesium has a large atomic radius, whereas Flourine has a very small one. This results in some atoms having a larger atomic radius than others. The more electrons there are (and the more shells that exist between the nucleus and valence shell), the harder it is for this goal to be accomplished. The protons in the nucleus of an atom are positively charged, and want to draw in the electrons surrounding it. Thus, in the d-block you don't have a clear periodic trend for electronegativity as you do in the s and p blocks.Īs in, an atom wanting more to hog electrons? (I'm assuming this because this comment is on the electronegativity video) In the d-block, you have many competing factors such as then need to have a full s and p, to energetic nuances due to the d subshell being partially full, and all these affect the electronegativity. That is why electronegativity goes down as you go down a Group of the periodic table in the s and p blocks (the d block is different). The more electrons the element has, the less advantage there is to gaining an electron compared to other members of the same Group. One has to do with how many electrons (not valence electrons, all of the electrons) the element has.

However, there are also competing factors. This need for being in an energetically favorable state gives rise to electronegativity. Thus, the closer an element is to completing both its outermost p and s subshells, the more energetically favorable it is to gain an electron to help that process along. I can give you some overly simplified basics though.Ītoms are particularly energetically stable if the s and p subshells of their outermost electron shell are both completely full. That is a complex issue and a bit of an advanced topic.