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Ionization energy trends

Definition of ion and ionization energy, and trends in ionization energy across a period and down a group.

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  • duskpin sapling style avatar for user Harshita kanal
    Where do the names-Cation and Anion come from?
    (12 votes)
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  • piceratops seed style avatar for user Love Oko
    why do the dips occur in the general trend from alkali metals to noble gases?
    (14 votes)
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    • duskpin ultimate style avatar for user Igor
      There are couple of reasons for that. One is that when electrons start to fill p orbital the ionization energy goes down a little. Another is when each of 3 p orbitals have one electron they start to pair as new ones are added (like when moving from nitrogen to oxygen).
      Check out this video for more details: https://www.youtube.com/watch?v=2AFPfg0Como
      (The exact answer starts at about five minutes from the beginning)
      (7 votes)
  • aqualine ultimate style avatar for user Biological
    Ionization energy only serves to remove an electron? It can't serve to add, in order to make anions?
    (5 votes)
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  • aqualine ultimate style avatar for user quintus.marco.polo
    Dumb question, but why do protons and electrons have charge?what is the difference between a neutron and a proton other that charge?
    (10 votes)
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  • male robot hal style avatar for user GeniusKid88
    How come mercury has a high ionization energy?
    (6 votes)
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    • duskpin seed style avatar for user Hrudya
      Mercury has the electronic configuration of
      [Xe] 4f(14) 5d(10) 6s(2).
      So you can see that like Beryllium, it is also have a completely filled s orbital. Also an atom is more stable if the orbitals are half filled or completely filled. So it is reluctant to loose the electron. Hence Mercury has a high ionization energy.
      (2 votes)
  • piceratops sapling style avatar for user haekele
    At the end of the video it was told that Radon has a lower ionization energy than hydrogen. How come this is possible, I thought that noble gases always are harder to ionize.
    (5 votes)
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    • starky seedling style avatar for user 1324354657
      You’re right that the noble gases have high ionization energy. However, they only have the highest ionization energy in their period because they are the most stable elements. Moving down the noble gases group, the ionization energy decreases. So while you are right that in general, noble gases have high ionization energy, it is still easier to take an electron away from some noble gases than others.
      (3 votes)
  • duskpin seed style avatar for user Yashonandini Bhargava
    If ionization energy is used to remove electrons, which energy is used to add electrons?
    (3 votes)
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  • blobby green style avatar for user uditdagar2002
    You said that IE decreases down the group because shell increases.
    But, protons also increase simultaneously in the nucleus which eventually increases force of attraction.
    Shouldn't the IE be higher down the group?
    (3 votes)
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  • male robot hal style avatar for user Alik Itkin
    Which one of the following needs more ionization evergy the element tc, or ir?
    (3 votes)
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    • piceratops ultimate style avatar for user Nicolas Posunko
      You need approximately 25% more energy to ionize a set number of Iridium atoms than you need to ionize the same number of Technetium atoms (NB: That's if we're talking about separating only the first electron from each atom, kind of about the first stage of ionisation). BTW, you can look up the information on atoms of other elements on Wikipedia—just enter the name of the element (or its symbol) and find "Ionization Energies" in the table of its properties on the right (or control-F search "Ionization Energies"). In this case, you would find the following data:

      1-st ionization energy for Tc — 702 kJ·mol−1
      2-nd ionization energy for Tc — 1470 kJ·mol−1
      3-rd ionization energy for Tc — 2850 kJ·mol−1

      1-st ionization energy for Ir — 880 kJ·mol−1
      2-nd ionization energy for Ir — 1600 kJ·mol−1
      (3 votes)
  • blobby green style avatar for user nathalie Mangulabnan
    what do you mean by "if you take an electron away from them then their outer shell will going to have electron configuration of the noble gas before it."?
    (2 votes)
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    • starky seed style avatar for user melissa
      removing that single electron in the outer shell (valence electrons are the outermost electrons) its electron configuration will resemble that of a noble gas (that has octet).
      EX- NA (Sodium) has one valence electron. By losing this electron it will have ten electrons (like NEON) and since in nature, "8 is the golden number" it will have both its "s" and "p" block full, without any outliers. Hope this help :)
      (2 votes)

Video transcript

- [Voiceover] So, let's talk a little bit about a word you might have heard and that is Ion. Let's talk about what it is and then we'll talk about trends in the periodic table on, on I guess how hard it is to make something an Ion. In particular how hard it is to make something a positive ion. So, an ion is just an atom or a molecule that has charge and it'll have charge if the protons are not equal to the electrons. Neutrons are obviously also constituent of atoms but neutrons are neutral. What you're gonna get your charge from are your protons or electrons. So, you're going to have a net charge. If your number of, number of protons, and this is for an atom or molecule. A molecule's just a bunch of, a bunch of atoms bonded together. If the number of protons does not equal the number of electrons. And you can have positive ions if the protons are more than the number of electrons, protons, or positive electrons or negative. And you can have negative ions if the number of electrons are greater than the number of protons. For example, for example, if you just had Hydrogen in it's neutral state has one proton and one electron, but if you were to take one of those electrons away then Hydrogen would have a positive charge and essentially it would just be, in its most common isotope it would just be a proton by itself. And so, when we talk about a positive ion like this where our protons are more than our electrons, the number of protons are more than the number of electrons, we call these cations, cations. Cation, once again, just another word positive ion. Likewise, we can have negative ions. So, say for example, Fluorine. So, Fluorine gains an electron, it's going to have a negative charge. It's gonna have a negative charge of negative one, and a negative ion we call an anion. And the way that I remember this is a kind of means the opposite or the negation of something. So, this is a negative ion. We've negating, you can somehow think we are negating the ion. So, with that out of the way, let's think about how hard it will be ionize different elements in the periodic table. In particular, how hard it is to turn them into cations. And to think about that, we'll introduce an idea called ionization energy. Ionization... Ionization energy... Energy... And this is defined, this is defined as the energy required, energy required to remove an electron, to remove an electron. So, it could've even been called cationization energy because you really see energy required to remove an electron and make the overall atom more positive. So, let's think about the trends. And we already have a little bit of background on the different groups of the periodic table. So, for example, if we were to focus on, especially we could look at group one, and we've already talked about how Hydrogen's a bit of a special case in group one but if we look at everything below Hydrogen. If we look at the Alkali, if we look at the Alkali metals here we've already talked about the fact that these are very willing to lose an electron. Why? Because if they lose an electron they get to the electron configuration of the noble gas before it. So, if Lithium loses an electron then it has an outer shell electron configuration of Helium. It has two outer electrons and that's kind of, we typically talk about the Octet Rule but if we're talking about characters like Lithium or Helium they're happy with two 'cause you can only put two electrons in that first shell. But all the rest of 'em, Sodium, Potassium, etc., etc., if you take an electron away from them then their outermost shell, well, all of them in their outermost shell they're going to have the electron configuration of the noble gas before it and for Sodium on down that outer shell is going to have that perfect eight. Lithium, if you remove an electron, it would get to Helium and it would have two electrons in its outer shell. So, you can imagine that the ionization energy right over here, the energy required to remove electrons from your Alkali Metals is very low. So, let me just write down this is... So, when I say low, I'm talking about low ionization energy. Low. Now, what happens as we move to the right of the periodic table? In fact, let's go all the way to the right on the periodic table. Well, if we go here to the Noble Gases, the Noble Gases we've already talked about. They're very, very, very stable. They don't want no one, they don't want their electron configurations messed with. So, it would be very hard... Neon on down has their eight electrons that (mumbling) Octet Rule. Helium has two which is full for the first shell, and so it's very hard to remove an electron from here, and so it has a very high ionization energy. Low energy, easy to remove electrons. Or especially the first electron, and then here you have a high ionization energy. I know you have trouble seeing that H. So, this is high, high ionization energy, and that's the general trend across the periodic table. As you go from left to right, you go from low ionization energy to high ionization energy. Now, what about trends up and down the periodic table? Well, within any group, if we, even if we look at the Alkali, if we look at the Alkali Metals right over here, if we're down at the bottom, if we're looking at, if we're looking at, say, Cesium right over here, that electron in the, one, two, three, four, five, six, in the sixth shell, that's going to be further from that one electron that Lithium has and its second shell. So, it's going to be, it's going to be further away. It's not going to be as closely bound to the nucleus, I guess you could say. So, this is going to be even, that one electron's gonna even easier to remove than the one electron in the outermost shell of Lithium. So, this one has even lower, even lower, even lower... And that's even going to be true of the Noble Gases out here that Xenon, that it's electrons in its outermost shell, even though it has eight valence electrons, they're further away from the nucleus, and so they're a little, the energy required to remove them is still going to be high but it's going to be lower than the energy from, from say Neon or Helium. So, this is low. So, once again, ionization energy low to high as we go from left to right, and low to high as we go from bottom to top. Or we could say a general trend that if we go from the bottom left to the top right we go from low ionization energy, very easy to remove an electron from these characters right over here to high ionization energy, very hard to move, remove an electron from these characters over here. And you can see it if, you could see in a trend of actual measured ionization energies and I like to see charts like this because it kind of show you where the periodic table came from when people noticed these kind of periodic trends. It's like, hey, it looks like there's some common patterns here. But on this one in particular we see on this axis we have ionization energy and electron volts, that's actually, it's literally a, this is units of energy. You could convert it to Joules if you like. Then over here, we're increasing the atomic numbers. So, we're (mumbling), we're starting with Hydrogen then we go to Helium, and we keep, and then we go, go from Hydrogen to Helium to Lithium and let me show you what's happening right over here. So, you go to Hydrogen to Helium. So, Helium here is very stable, so it's very hard to remove an electron. And then you go to Lithium. Lithium, as we said, this is an Alkali metal. You remove an electron, it gets to a stable state. So, it takes very low energy to remove that electron. And then as we go from left to right on the periodic table, as we go from Alkali Metal to Noble Gases we see that the ionization energy increases. And there are these little dips here which you could think about why these... (mumbling) Or theorize why these dips are occurring, what you see in this general trend as we go from Alkali Metals to Noble Gases. Alkali Metals to Noble Gases. Alkali Metals to Noble Gases. Now, one thing you might be saying is, "Hey, look, you had from here to here, "that's the same distance as here to here, "but now we have a larger distance here. "What's going on here?" Well, we have to remember now we have all of our D block elements. So, now, once we get, once we get to the, once we get over here we're now adding all of the D block elements. (mumbling) On the fourth period and so we have those, we have those added here, so you have D block elements, D block elements, and then here you have you F and D block elements. And so, you see the general trend that your Alkali, your Alkali Metals are very low ionization energy. Your Noble Gases, very high ionization energy. But as they get, as the atoms get larger and larger the ionization energy goes lower and lower, and sends something like Radon, which even though it's Noble Gas it's ionization energy because those outermost electrons are further away from the nucleus or they're quite far away from the nucleus, that its ionization energy is actually, its ionization energy is actually less then that of Hydrogen. Anyway, hope you found that interesting.