Within a period, the IE 1 generally increases with increasing Z. Down a group, the IE 1 value generally decreases with increasing Z. There are some systematic deviations from this trend, however. Note that the ionization energy of boron atomic number 5 is less than that of beryllium atomic number 4 even though the nuclear charge of boron is greater by one proton.
This can be explained because the energy of the subshells increases as l increases, due to penetration and shielding. Within any one shell, the s electrons are lower in energy than the p electrons. This means that an s electron is harder to remove from an atom than a p electron in the same shell. The electron removed during the ionization of beryllium [He]2 s 2 is an s electron, whereas the electron removed during the ionization of boron [He]2 s 2 2 p 1 is a p electron; this results in lower first ionization energy for boron, even though its nuclear charge is greater by one proton.
Thus, we see a small deviation from the predicted trend occurring each time a new subshell begins. Another deviation occurs as orbitals become more than one-half filled.
The first ionization energy for oxygen is slightly less than that for nitrogen, despite the trend in increasing IE 1 values across a period. For oxygen, removing one electron will eliminate the electron-electron repulsion caused by pairing the electrons in the 2 p orbital and will result in a half-filled orbital which is energetically favorable.
Analogous changes occur in succeeding periods. Removing an electron from a cation is more difficult than removing an electron from a neutral atom because of the greater electrostatic attraction to the cation. Likewise, removing an electron from a cation with a higher positive charge is more difficult than removing an electron from an ion with a lower charge.
Thus, successive ionization energies for one element always increase. As seen in Table 1, there is a large increase in the ionization energies for each element. This jump corresponds to the removal of the core electrons, which are harder to remove than the valence electrons. For example, Sc and Ga both have three valence electrons, so the rapid increase in ionization energy occurs after the third ionization. This text is adapted from OpenStax Chemistry 2e, Section 6.
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Chapter 4: Chemical Quantities and Aqueous Reactions. Chapter 5: Gases. Chapter 6: Thermochemistry. Chapter 7: Electronic Structure of Atoms. Chapter 9: Chemical Bonding: Basic Concepts. Chapter Liquids, Solids, and Intermolecular Forces. Large atoms or molecules have low ionization energy, while small molecules tend to have higher ionization energies.
The ionization energy is different for electrons of different atomic or molecular orbitals. More generally, the nth ionization energy is the energy required to strip off the nth electron after the first n-1 electrons have been removed.
It is considered a measure of the tendency of an atom or ion to surrender an electron or the strength of the electron binding. The greater the ionization energy, the more difficult it is to remove an electron. The ionization energy may be an indicator of the reactivity of an element.
Elements with a low ionization energy tend to be reducing agents and form cations, which in turn combine with anions to form salts. Moving left to right within a period or upward within a group, the first ionization energy generally increases. As the atomic radius decreases, it becomes harder to remove an electron that is closer to a more positively charged nucleus.
Conversely, as one progresses down a group on the periodic table, the ionization energy will likely decrease since the valence electrons are farther away from the nucleus and experience greater shielding.
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. 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. 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. 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.
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