- Periodic trends and Coulomb's law
- Atomic and ionic radii
- Ionization energy: group trend
- Ionization energy: period trend
- First and second ionization energy
- Worked example: Identifying an element from successive ionization energies
- Electron affinity: period trend
- Periodic trends
Ionization energy: period trend
An element's first ionization energy is the energy required to remove the outermost, or least bound, electron from a neutral atom of the element. On the periodic table, first ionization energy generally increases as you move left to right across a period. This is due to increasing nuclear charge, which results in the outermost electron being more strongly bound to the nucleus. Created by Jay.
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- I still dont really get it about the decreasing of IE from N to O. Can give me a simpler explanation?(28 votes)
- You know the maximum capacity of p orbital is 6 electrons. so N has three 2p electrons and O has four 2p electrons. This case is kind of exception. Since Nitrogen is half filled coz it has 3 electrons out of 6 electrons that can be occupied in p orbital it has higher ionization energy compared to that of oxygen.
Half filled and full filled orbtials usually tend to show higher ionization energy . Thus the noble gas elements have high ionization energy. (FYI Fluorine has the highest ionization energy)(11 votes)
- If I'm not mistaken, the electrons (electron clouds or whatever it is called) are in constant motion around an atom. If this is the case how can the inner shell electrons "shield" them constantly? Wouldn't the ionization energy change?(10 votes)
- Although we can describe electrons in terms of mathematics, actually knowing what they look like is beyond us. You have probably heard of wave-particle duality which describes the behaviour of sub-atomic particles, such as electrons - they have the properties of being both waves and particles at the same time. If you wish to think of electrons as particles, then you have to accept they are everywhere at once, which is down to their wave function properties - they are not moving around from one place to another. If they are everywhere at once, then they are providing constant shielding.
This is quantum mechanics which is extremely weird and can be very counter-intuitive.(14 votes)
- Do electrons move between shells? If the orbitals are simply probability zones, would those zones ever overlap? Are the numbers of electrons per orbital also just be an average or a probability?(5 votes)
- 1. Electrons jump up to higher shells when excited and fall back down and release that energy as photons.
2. Yes, orbitals can overlap.
3. Not quite sure what you mean. Orbitals are probability locations of where a certain number of electrons are. We aren't uncertain about the umber of electrons, but their location and momentum. So I guess yes, if we are uncertain about whether there are 2 electrons in an s orbital, that could mean that there are fewer or more in there at a point in time.(5 votes)
- What does it mean by shielding?(3 votes)
- Shielding is when the inner electrons partially block the attractive force between the positively charged nucleus and the outer electron(s). This makes it easier to remove the outer electron(s) from the atom.(4 votes)
- What exactly is Effective Nuclear Charge?(3 votes)
- It’s the charge of the nucleus that an outer electron feels. Core electrons shield the outer electrons from some of the positive charge of the nucleus.(4 votes)
- So, is ionization energy the energy needed to remove an electron from an atom? I'm confused. I thought ionization energy was to remove all the outer electrons.(3 votes)
- ionization energy is defined as the energy required to remove the most loosely bound(Valence electron) electron from an isolated atom in a gaseous state.(1 vote)
- Can anyone explain me why does the Ionisation energy in a second period decreases from Be ( 900 KJ/Aton ) to B (800KJ/Atom) but the trend clearly states that in a period the I.E increases from left to right.(4 votes)
- First, it is only a trend, not a strict law of nature. There are many exceptions to the trend.
Careful examination will show that there is a sudden drop in ionization energy for members of the Boron Group (Group 13) compared to the previous element. That is because Group 13 is where the elements switch from filling s or back-filling d orbitals to filling p orbitals. The p orbitals are farther from the nucleus and thus are not as strongly held. But as you continue filling the p subshell, the ionization energy goes back up because the nuclear charge is increasing (amongst other reasons) while you are filling the same subshell.
You see the same pattern when you move from Group 18 to the next element which is the Group 1 element of the next period. You have gone from filling the p subshell to filling the s subshell of the next higher shell, so it is much less tightly held because the new s subshell is considerably farther away from the nucleus.(1 vote)
- my question is that why is there such a big jump in ionization energy from Li to Be(2 votes)
- 1. When an electron is removed from Li, all of its remaining electrons are in the n=1 shell. That is, losing an electron puts it in a particularly stable state, so it is relatively easy to remove this electron.
2. Be has a higher effective nuclear charge than Li, since it has more protons but the same number of core/inner electrons. Thus, the valence electrons are held more strongly and so it is harder to remove them from Be than Li.(3 votes)
- Isn't the atomic radius of beryllium smaller than the atomic radius of lithium? (there are more protons and still the same amount of shielding so the electrons are a bit closer to the nucleus) Shouldn't that mean that the electrons are closer to the nucleus in beryllium instead of the same distance?(2 votes)
- The electrons in Beryllium are indeed 'closer' to the nucleus. That is also the reason Ionization energy of beryllium is more than that of Lithium.(2 votes)
- The noble gas xenon can form many compounds, but neon,
another noble gas, doesn't form compounds. why is that so?Why
couldn't Ne form NeF4 like XeF4? please answer(1 vote)
- Have a look at a periodic table (e.g. https://ptable.com) — what do you notice about Xe vs. Ne?
Elements lower on the table are larger and the outer electrons are less strongly attracted to the nucleus. This is partly because they are further away, but also because there are more "layers" (shells) of electrons shielding them from the nucleus.
What do you think that would mean about how easy it is for an atom to share outer (valence) electrons?
Does that help you answer your question?(4 votes)
- [Instructor] In this video, let's look at the periodic trends for ionization energy. So, for this period, as we go across from lithium, all the way over to neon, so as we go this way, across our periodic table, we can see, in general, there's an increase in the ionization energy. So, lithium is positive 520 kilojoules per mole. Beryllium's goes up to 900 kilojoules per mole, and then again, in general, we see this increase in ionization energies going over to neon. So, going across a period, there's an increase in the ionization energy. And that's because, as we go across our period, there's an increase in the effective nuclear charge. So, increase in Z effective. And remember, the formula for that is the effective nuclear charge is equal to the actual number of protons, which is Z, and from that we subtract S, which is the average number of inner electrons shielding our outer electrons. So, let's examine this in more detail, looking at lithium and beryllium. Lithium has atomic number three, so three protons in the nucleus, so positive three charge, and lithium's electron configuration we know is one s two, two s one. So, two electrons in our one s orbital, and one electron in the two s orbital. Beryllium has one more proton and one more electron. So one more proton in the nucleus, so a plus 4 charge, and for beryllium, the electron configuration is one s two, two s two. So two electrons in the one s orbital, and then two electrons in the two s orbital. Let's calculate the effective nuclear charge for both of these, and first, we'll start with lithium. So for lithium, lithium has a plus three charge in the nucleus, so the effective nuclear charge is equal to positive three, and from that we subtract the average number of inner electrons shielding our outer electrons, in this case, we have these two inner, or core electrons, that are shielding our outer electron, our valence electron, from this full positive three charge. So we know that like charges repel, so this electron is going to repel this electron a little bit, and this electron repels this electron. And these two inner core electrons of lithium have a shielding effect, they protect the outer electron from the full positive three charge. So there's two shielding electrons, so for a quick effective nuclear charge calculation positive three minus two gives us a value of plus one for the effective nuclear charge. So, it's like this outer electron of lithium is feeling a nuclear charge of plus one, which pulls it toward the nucleus, right? So, there's an attractive force between the outer electron and our nucleus. Now, the actual calculation for this um, Z is-- S I should say, does not have to be an integer, and the actual value for lithium is approximately one point three, but our quick, crude calculation tells us positive one. Let's do the same calculation for beryllium, so the effective nuclear charge for beryllium is equal to the number of protons, right, which for beryllium is positive four, and from that, we subtract the number of inner electrons that are shielding the outer electrons. So, it's a similar situation, we have two inner electrons that are shielding this outer electron, they're repelling this outer electron, shielding the outer electron from the full positive four charge of the nucleus. SO we say there are two inner electrons, so the effective nuclear charge is positive four minus two, giving us an effective nuclear charge of positive two. In reality, the effective nuclear charge is approximately one point nine, and that's because beryllium has another electron in its two s orbital over here, which does effect this electron a little bit. It repels it a little bit, and so it actually deceases the effective nuclear charge to about, one point nine. But again, for a quick calculation, positive two works. So, the outer electron for beryllium, let's just choose this one again, is feeling an effective nuclear charge of positive two, which means that, it's going to be pulled closer to the nucleus, there's a greater attractive force on this outer electron for beryllium, as compared to this outer electron for lithium. The effective nuclear charge is only plus one for this outer electron, and because of this, the beryllium atom is smaller, right? The two s orbital gets smaller, and the atom itself is smaller. Beryllium is smaller than lithium. So this outer electron here, let me switch colors again, this outer electron for beryllium is closer to the nucleus than the outer electron for lithium. It feels a greater attractive force, and therefore it takes more energy to pull this electron away from the neutral beryllium atom, and that's the reason for the higher ionization energy. So beryllium has an ionization energy of positive 900 kilojoules per mole, compared to lithium's of 520 kilojoules per mole. So it has to do with the effective nuclear charge. So far we've compared lithium and beryllium and we saw that the ionization energy went from positive 520 kilojoules per mole to 900 kilojoules per mole, and we said that was because of the increased effective nuclear charge for beryllium, but as we go from beryllium to boron, there's still an increased effective nuclear charge, but notice our ionization energy goes from 900 kilojoules per mole for beryllium to only 800 kilojoules per mole for boron, so there's a slight decrease in the ionization energy. And let's look at the electron configuration of boron to see if we can explain that. Boron has five electrons, so the electron configuration is one s two, two s two, and two p one. So that fifth electron goes into a two p orbital, and the two p orbital is higher in energy than a two s orbital, which means the electron in the two p orbital is on average, further away from the nucleus that the two electron in the two s orbital. So if we just sketch this out really quickly, let's say that's my two s orbital, I have two electrons in there, and this one electron in the two p orbital is on average further away from the nucleus. So, those two electrons in the two s orbital actually can repel this electron in the two p orbital. So, there's a little bit extra shielding there of the two p electron from the full attraction of the nucleus, right? So, even though we have five protons in the nucleus, and a positive five charge for boron, the fact that these two s electrons add a little bit of extra shielding means it's easier to pull this electron away. So, it turns out to be a little bit easier to pull this electron in the two p orbital away due to these two s electrons. And that's the reason for this slight decrease in ionization energy. As we go from boron to carbon, we see an increase in ionization energy, from carbon to nitrogen, an increase in ionization energy. Again, we attribute that to increased effective nuclear charge, but when we go from nitrogen to oxygen, we see a slight decrease again. From about 1400 kilojoules per mole, down to about 1300 kilojoules per mole for oxygen. So, let's see if we can explain that by writing out some electron configurations for nitrogen and oxygen. Nitrogen has seven electrons to think about. So it's electron configuration is one s two, two s two, and two p three. So that takes care of all seven electrons. For oxygen, we have another electron, so one s two, two s two, two p four is the electron configuration for oxygen. Let's just draw using orbital notation the two s orbital and the two p orbital. So for nitrogen, here's our two s orbital. We have two electrons in there, so let's draw in our two electrons. And for our two p orbitals, we have three electrons. So here are the two p orbitals, and let's draw in our three electrons using orbital notation. Let's do the same thing for oxygen. So there's the two s orbital for oxygen, which is full, so we'll sketch in those two electrons, and we have four electrons in the two p orbitals. So let me draw in the two p orbitals. There's one electron, there's two, there's three, and notice what happens when we add the fourth electron. We're adding it to an orbital that already has an electron in it, so when I add that fourth electron to the two p orbital, it's repelled by the electron that's already there, which means it's easier to remove one of those electrons, so electrons have like charges, and like charges repel. And so that's the reason for this slight decrease in ionization energy. So, it turns out to be a little bit easier to remove an electron from an oxygen atom, than nitrogen, due to this repulsion in this two p orbital. From there on, we see our general trend again. The ionization energy for fluorine is up to 1681, and then again for neon, we see an increase in the ionization energy due to the increased effective nuclear charge.