This question has been bugging me for quite some time now...what is the difference between core charge and electronegativity?
I know that core charge is the attraction between the nucleus and the valence electrons while electronegativity is the ability of an atom to attract electrons towards itself, but in terms of their trends in the periodic table, I'm not sure what the difference is.
For example, apparently core charge remains the same down a group but how can that be? Shouldn't it decrease since there are more shells, meaning the valence electron is further from the nucleus, so there is less attraction between the valence electrons and the nucleus?
Another example is how, in the explanations of why electronegativity decreases down a group and increases across a period, the textbook mentions attraction between the valence electrons and the nucleus, and not between the atom and other electrons, like the definition of electronegativity states.
Also I'm not sure of the difference between the reason why electronegativity increases across a period and why core charge increases across a period.
I am so confused by this. Any help is appreciated! Thanks 
I haven't learnt much about core charge, but:
- it increases across a period because the number of valence electrons increases, whilst the protons also increase, hence the attractive force increases when you move across the table (more protons + electrons = stronger attraction).
- As the number of shells increases, the core charge decreases? Since as your number of shells increases, the attractive force between the inner shells and the nucleus decreases.
I reckon its similar to electronegativity. Theoretically it should increase as electronegativity increases, but I don't think that's the case (after googling).
-Idk what you mean by electronegativity states, but it's because electronegativity also highlights how badly an atom wants to keep its own electrons. If its strong enough to attract electrons from another element, it must be strong enough to keep its own from going to that element? So if it wants to keep its electrons, the attractive force between the nucleus and the valence electrons must be stronger. This can be achieved by: more electrons in the outer shell, and/or fewer shells. Hence fluorine...
The elements with high electronegativities tend to hog the electrons, i.e. they keep their own, and take others. So basically, both the relationship between the atom (nucleus and valence electrons) and other elements are kinda linked.
Finally, since core charge is similar to electronegativity, pretty sure it changes similarly. Both increase with ability of the nucleus to attract electrons, hence why across the period, you have more electrons, meaning a stronger attractive force!
Another question - with energy profile diagrams, does the position of the y intercept matter?
Like does it matter if you say its 100 kJ or 0 kJ - coz what you're stating in that graph is the activation energy and the delta H value - so as long as the magnitudes are right is your graph right?
Is this correct?
It's usually measured from 0. If the reactants are at 100 kJ, my y-int would be at 100 kJ. I wouldn't start at 0 personally and I'm not entirely sure if you're allowed to. This would probs be a question for school teachers, since its risky.
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