Is there a difference between ionization energy and ionization potential?

In summary, Wikipedia states that ionization energy and ionization potential are synonymous terms. Ionization energy is the amount of energy needed to remove an electron from an atom, and a molecule with a low ionization energy can be more easily ionized. However, it may seem counterintuitive to call this a "low ionization potential" since it implies that it is not easy to ionize the molecule. This raises the question of whether ionization energy and ionization potential are the same thing or opposites. Additionally, the explanation for why ionization energy decreases down a group is that the electrons are further from the nucleus, but the effective nuclear charge increases. This is due to the shielding effect of core electrons. However, the relationship between
  • #1
LogicX
181
1
Wikipedia says they are synonymous.

Ionization energy is how much energy it takes to abstract an electron from an atom. A molecule with a low ionization energy can more easily be ionized. It seems weird to call this a low ionization potential though. You are trying to say it can easily be ionized, but you would call that a "low ionization potential"? That makes it sound like it ISN'T easy to ionize it.

So, are they the same thing or are they opposites?

Also, why does IE decrease down a group? The explanation given is that the electrons are further from the nucleus. But Z(effective) increases down a group. Who cares if they are further away if they are feeling more charge from the nucleus?
 
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  • #2
Yes, but as you move down a group, the valence shell is now "shielded" in part by the electrons in the non-valence shells as well.
 
  • #3
daveb said:
Yes, but as you move down a group, the valence shell is now "shielded" in part by the electrons in the non-valence shells as well.

I know there are core electrons that shield. But as you move down the core e- are not as effective at shielding due to increased core shell size. The effect is that Z* increases slightly down a group. My question is how can the electromagnetic force on an electron at that point be larger but it is also easier to pull that electron away?
 
  • #4
Ionization energy and ionization potential are used interchangeably in my experience. If you want to argue for differentiating between the two, be my guest, but chemists can be awfully slow in adopting new terminology standards (I still call ethene "ethylene," after all).

LogicX said:
I know there are core electrons that shield. But as you move down the core e- are not as effective at shielding due to increased core shell size. The effect is that Z* increases slightly down a group. My question is how can the electromagnetic force on an electron at that point be larger but it is also easier to pull that electron away?

I presume that you've seen the spherical harmonics plotted for the higher-energy shells. (They're on Wikipedia somewhere, in case you haven't - I remember seeing them not too long ago.) It's not just that they're farther from the nucleus in a general sense - the probability of finding an electron is noticeably smaller nearer the nucleus for, say, an electron in the sixth energy level, relative to the second energy level.

There's probably also some benefit to reducing electron-electron repulsion by knocking out an electron further down the groups, although that's kind of a handwavy thing to say, and I can never remember just how well it holds for heavier elements in all cases.
 
  • #5
Mike H said:
I presume that you've seen the spherical harmonics plotted for the higher-energy shells. (They're on Wikipedia somewhere, in case you haven't - I remember seeing them not too long ago.) It's not just that they're farther from the nucleus in a general sense - the probability of finding an electron is noticeably smaller nearer the nucleus for, say, an electron in the sixth energy level, relative to the second energy level.

I understand this but I think the problem still stands. Trends in ionization energy should be directly related to Z*. It makes sense to do it this way, and it is how the trend across a period is explained (Z* inc. across a period, so it is harder to pull an electron off). I don't see how any sort of distance relation should matter.
 
  • #6
Look at it in the classical E&M sense. The force exerted on a charged particle from another charged particle is directly proportional to the product of the charges and inversely proportional to the square of the distance between those charges, so if distance increases, then the force exerted is less, and the electron is less bound to the nucleus.
 
  • #7
daveb said:
Look at it in the classical E&M sense. The force exerted on a charged particle from another charged particle is directly proportional to the product of the charges and inversely proportional to the square of the distance between those charges, so if distance increases, then the force exerted is less, and the electron is less bound to the nucleus.

So Z* at one distance away has a different force exerted on it than the same Z* at a different distance?

Ah ok, for some reason I was thinking that it just mattered how much charge it "felt", not how much it felt and how far away it was.
 
  • #8
LogicX said:
So Z* at one distance away has a different force exerted on it than the same Z* at a different distance?

Ah ok, for some reason I was thinking that it just mattered how much charge it "felt", not how much it felt and how far away it was.

As daveb mentioned, it's Coulomb's law all over again. The force a charge "feels" is explicitly dependent on the square of the distance. If you have a distance of 2 length units between charges, you'd divide the product of the charges by 4 (=22). If you have a distance of 4 units, you'd divide the product by 16 (=42).

To take a simplified example (numbers from WebElements) = Z* for Be's valence electron is 1.91, while the valence shell orbital radius is 2.05 AU. For Ba, Z* is 7.6 and the radius is 4.45 AU. Let the product of the electron charge and the coefficient in Coulomb's law equal x. Do this simplified calculation given Coulomb's law, and for Be, you get ~ 0.45x, while for Ba you get ~ 0.38x - that's a 15% difference. The gap in their first ionization energies is greater than this, but chalk that up to a purely classical treatment.

(I know I played a bit fast and loose with units in this post. No need to flagellate me over it.)
 
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1. What is ionization energy and ionization potential?

Ionization energy and ionization potential both refer to the amount of energy required to remove an electron from an atom or molecule, resulting in the formation of an ion. They are essentially the same concept, but ionization potential is more commonly used in atomic physics while ionization energy is more commonly used in chemistry.

2. Is there a difference in the units used to measure ionization energy and ionization potential?

No, both ionization energy and ionization potential are typically measured in units of energy, such as joules (J) or electron volts (eV). However, some sources may use different units, such as kilojoules per mole (kJ/mol) for ionization energy.

3. Can ionization energy and ionization potential vary for different atoms or molecules?

Yes, the amount of energy required to remove an electron from an atom or molecule can vary depending on its atomic or molecular structure. This is because the number of protons, the distance between the nucleus and the outermost electron, and other factors can affect the strength of the electron's attraction to the nucleus.

4. What is the relationship between ionization energy and ionization potential?

As stated earlier, ionization energy and ionization potential are essentially the same concept. However, ionization potential is often used to refer to the energy required to remove an electron from a gaseous atom, while ionization energy can refer to the energy required to remove an electron from an atom in any state (solid, liquid, or gas).

5. How are ionization energy and ionization potential measured?

Ionization energy and ionization potential can be measured experimentally using techniques such as photoelectron spectroscopy or mass spectrometry. These methods involve bombarding atoms or molecules with high-energy particles or photons and measuring the energy required to remove an electron. Theoretical calculations can also be used to estimate ionization energy and ionization potential values.

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