- Periodic trends
- Atomic radius trends on periodic table
- Atomic and ionic radii
- Mini-video on ion size
- Ionization energy trends
- Ionization energy: period trend
- First and second ionization energy
- Electron affinity: period trend
- Electronegativity and bonding
- Metallic nature
- Periodic trends and Coulomb's law
- Worked example: Identifying an element from successive ionization energies
- Ionization energy: group trend
Definition of ion and ionization energy, and trends in ionization energy across a period and down a group.
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- Where do the names-Cation and Anion come from?(12 votes)
- The etymology of Cat- and An- is down and up respectively. Cations are positively charged and anions are negatively charged. Cathodes attract cations and anodes attract anions.(9 votes)
- why do the dips occur in the general trend from alkali metals to noble gases?(14 votes)
- 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)
- 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)
- There are many things that make up the particles in an atom, including Quarks (there are 6 "flavors", up, down, top, bottom, strange, and charm), and leptons. The most common lepton is an electron. Quarks also make up neutrons along with protons. Different "flavors" of quarks in different combinations can come out as protons or neutrons. Honestly, if you don't think this is enough just go to the link in the above answer or come here.
- Ionization energy only serves to remove an electron? It can't serve to add, in order to make anions?(4 votes)
- The energy that is involved when you add an electron to an atom is called the electron affinity.(16 votes)
- How come mercury has a high ionization energy?(5 votes)
- 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.(1 vote)
- 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.(4 votes)
- 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.(2 votes)
- 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)
- The evidence says that is not the case though. Going down a group the electrons get further and further away from the nucleus which means they are easier and easier to be ionised.(3 votes)
- In a question asking to explain the trend in 1st ionisation energy, is it sufficient to talk about how ionisation takes you either towards or further from the octet configuration (i.e. ionising Mg gets you closer to the electron configuration of Ne but ionising F gets you further away from the electron configuration of Ne) or should I also mention effective nuclear charge and shielding and whatnot. It's been a while since I've covered this and we're going over it again now. My teacher hasn't mentioned effective nuclear charge at all and a tutoring centre told me I should mention effective nuclear charge.(2 votes)
- Well it really depends on how thorough of an answer your teacher wants. The issue is that there are several reasons which influence an element's ionization energy. Stability because of a noble gas configuration will cause the ionization to be large, but having a greater effective nuclear charge will also generate a greater ionization energy.
But these explanations won't work for every situation. For example we would need to know about spin-pairing and the stability associated with half-filled subshells to explain why nitrogen has a higher first ionization energy compared to oxygen, even though oxygen has a greater effective nuclear charge.
So there is quite a lot of detail you could potentially go into with this question, but ultimately it depends on how much detail a teacher is seeking which we wouldn't know here.
Hope that helps.(4 votes)
- What is the ionization energy trends for the d and f block elements?(2 votes)
- [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.