Trend of First Ionization Energy in the Periodic Table

In summary, the atomic radius does not have a direct impact on ionization energy. Other factors such as the energy of the orbital the electron is on also play a significant role. Additionally, the concept of atomic radius is not well-defined and its relationship to other properties is not straightforward. Learning about orbitals and bond lengths may provide a better understanding of the topic.
  • #1
Itskitty
7
0
Why do metals generally have lower ionization energies than non-metals?

I mean, doesn't ionization energy depend on the atomic radius?

And the atomic radius is in turn dependent on the shell and the protons.

According to these factors, the atomic radius of Sodium should be smaller than
lower period non-metals such as Bromine.

Smaller atomic radius means that the proton can attract valence electrons much more strongly.

Therefore, Sodium's IE must be higher than that of Bromine.

But it's not the case. Why?
 
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  • #2
ok I'm wrong here about the trend

so we actually have to look at the trend in atomic radius of metals and non-metals separately.

Why?
 
  • #3
a clearer question is:
Why doesn't non-metals in the (n+1)th shell
have greater radius than metals in the nth shell
although it is the case for metals in the n+1 th shell?
 
  • #4
Itskitty said:
doesn't ionization energy depend on the atomic radius?

No, it doesn't. There is no such dependency, so all your thinking is based on a wrong assumption.

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  • #5
No?
But i thought tht smaller radius means tht the nucleus has a tighter hold on the valence electron
ths increasing the ionization energy?

Would u please correct my misunderstanding here?
 
  • #6
To some extent you are right, but this approach is way too simplified to give correct results. It would work for equivalent electrons. But electrons are not equivalent, they occupy orbitals - and amount of energy needed to remove electron depends on the energy of orbital the electron is on. This dominates the situation.
 
  • #7
I don't see it in my chemistry book (chang's general chemistry) could you tell me what topic of chemistry will help me to understand this?
Thanks so much!
 
  • #8
Any intro to quantum chemistry and quantum numbers will do. Don't have to be mathematically heavy.

--
 
  • #9
Borek: Well the first ionization potential does actually have one radius-related property, which is that there's a (not very well-)known relation to the asymptotic behavior of the electronic wave function (or density). (E.g. Katriel and Davidson, PNAS, v77, no 8, p4403)

I have to say I think those introductory/general chem textbooks fret far too much over atomic radii. As you said, there are no simple rules - and I'd add that it's not a terribly useful property. From a theoretical POV you'd be better learning more about orbitals, and from a practical/structural POV, one would be better off learning the lengths of common bonds.

So don't fret Itskitty, there's no absolute definition of what an atomic radius even is, much less any absolute rules on how it corresponds to other properties. I still find myself gaining new insights all the time, after spending years studying the stuff!
 

1. What is first ionization energy?

First ionization energy is the minimum amount of energy required to remove the outermost electron from an atom in its gaseous state.

2. How does first ionization energy change across the periodic table?

First ionization energy generally increases from left to right across a period. This is due to the increasing number of protons in the nucleus, which leads to a stronger attraction for the outermost electron.

3. Why does first ionization energy decrease down a group in the periodic table?

The outermost electrons in atoms in the same group have similar levels of energy, and as you move down the group, the distance between the outermost electron and the nucleus increases. This results in a weaker attraction between the two, making it easier to remove the outermost electron and decreasing the first ionization energy.

4. How does the trend of first ionization energy relate to the reactivity of elements?

Elements with low first ionization energy are more likely to lose electrons and become positively charged ions, making them more reactive. On the other hand, elements with high first ionization energy have a stronger hold on their electrons and are less likely to react.

5. Are there any exceptions to the trend of first ionization energy in the periodic table?

There are a few exceptions to the trend, particularly in the transition metals. For example, chromium and copper have lower first ionization energies than expected due to their half-filled and fully-filled subshells, respectively.

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