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 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. 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.
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. Let's look at lithium. So down here, we'll draw lithium. Lithium has an atomic number of three, so three protons in the nucleus. And in a neutral atom, three electrons. So the electron configuration is 1s2, 2s1. So there are two electrons in the first energy level and they're in an s orbital. So I'm going to go ahead and draw those in here.
So these two electrons I just drew represent the two electrons in the first energy level. In the second energy level, there's one more electron. So I'm going to put that electron down here like that. So for lithium, if we were to take an electron away, the one that's most likely to leave would be this outermost electron here, the one in the 2s orbital.
So if you apply kilojoules per mole of energy, you can pull away that electron. And so if you did that, you'd be left with a plus 3 charge in the nucleus. And you would still have your electrons in the 1s orbital, so I'm going to go ahead and draw those in there, but you've taken away that outer electron. And so therefore, you'd have a lithium cation here. You'd have Li plus 1, because you have three positive charges in the nucleus and only two electrons now. So 3 minus 2 gives you plus 1.
The electron configuration for the lithium cation would therefore be 1s2 because we pulled away that outer electron in the 2s orbital. So this is the picture for the ionization of hydrogen and lithium. And we're going to examine some of the factors that affect the ionization energy. And so first we'll talk about nuclear charge. So let me go ahead and write nuclear charge here. So the idea of nuclear charge is the more positive charges you have in your nucleus, the more of an attractive force the electron would feel.
And so therefore, the harder it would be to pull that electron away. So in general, you could think about increased nuclear charge. That would want to increase the ionization energy. Because again, there's a greater attractive force for the electrons. So let's look at these two situations, and let's think about hydrogen first. So hydrogen has a plus 1 charge in the nucleus.
And this one electron here would be pulled to the nucleus by that positive charge. If we look at lithium, plus 3 in the nucleus. So that's a greater nuclear charge. So just thinking about nuclear charge alone, you would think, oh, well this electron might be pulled in even more than with hydrogen, because plus 3 is greater than plus 1.
And so just thinking about nuclear charge for these two things, that would seem to indicate that lithium's outer electron would have a greater attractive force for the nucleus.
So therefore, you might think it might take more energy to pull that electron away. So just thinking about nuclear charge, we might think an increase in the ionization energy. Next, let's talk about electron shielding. So electron shielding, or you could also call it electronic screening.
So the idea of electron shielding is the inner shell electrons are going to shield the outer electrons from the positive charge of the nucleus. And let's look at lithium for an example of that. So we have these two inner shell electrons are going to repel the outer shell electrons.
So this electron in blue is going to repel this electron in green, and this electron in blue is going to repel this electron in green. And so they're going to shield that outer electron in green from that positive 3 charge, because electrons repel other electrons.
Like charges repel other like charges. And so that's the idea of electron shielding or electron screening. And so thinking about just this factor, for lithium, these two inner shell electrons are going to shield that outer shell electron. There's going to be a force in the opposite direction, if you will.
And so that means that it would be easier to take that outer electron away due to the repulsive force of those electrons. And so if we just think about electron shielding or electron screening by itself, it would be easier to take away lithium's outer electron due to the shielding effect. Diagram 4: showing decreasing pattern of electron affinities of elements from top to bottom 9. As indicated above, the elements to the right side of periodic table diagram 3 have tendency to receive the electron while the one at the left are more electropositive.
Also, from left to right, the metallic characteristics of elements decrease 4. The difference of electronegativity or ionization energies between two reacting elements determine the fate of the type of bond.
For example, there is a big difference of ionization energies and electronegativity between Na and. Therefore, sodium completely removes the electron from its outermost orbital and chlorine completely accepts the electron, and as a result we have an ionic bond 4.
However, in cases where there is no difference in electronegativity, the sharing of electrons produces a covalent bond. For example, electronegativity of Hydrogen is 2. What bond do they form when chemically combined?
According to periodic trends, one would assume that calcium, being to the left of gallium, would have the lower ionization energy. Explain, in terms of orbitals, why these numbers make sense. Calcium, however, has a fully stable 4s orbital as its valence orbital, which you would have to disrupt to take an electron away from.
Periodic Table and Trend of Ionization Energies As described above, ionization energies are dependent upon the atomic radius. Example of how ionization energy increases as succeeding electrons are taken away.
The Effects of Electron Shells on Ionization Energy Electron orbitals are separated into various shells which have strong impacts on the ionization energies of the various electrons. Ionization Energy and Electron Affinity--Similar Trend Both ionization energy and electron affinity have similar trend in the periodic table. B 27 C Prediction of Covalent and Ionic Bonds The difference of electronegativity or ionization energies between two reacting elements determine the fate of the type of bond.
Questions 1 By looking at following electronic configuration of elements, can you predict which element has the lowest first ionization energy?
Covalent Polar Covalent Ionic 8 Ionization energy, when supplied to an atom, results in a n Anion and a proton Cation and a proton Cation and an Electron Anion and an electron 9 Low first ionization energy is considered a property of Metals Nonmetals 10 Gallium has a first ionization energy of
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