Why is the ionization energy of Cl lower than F, but higher than that of S?

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Welcome back everyone. Germanium has a lower ionization energy than silicon but is found to have a higher ionization energy than gallium. We need to explain why. So in choice A it says that Germanium has a lower ionization energy than silicon because its shielding effect is stronger. But on the other hand, Germanium has a higher ionization energy than gallium because its effective nuclear charge is higher in choice B it states that Germanium has a lower ionization energy than silicon because its shielding effect is stronger. But Germanium has a higher ionization energy than gallium because its effective nuclear charge is lower. In choice C it says that Germanium has a lower ionization energy than silicon because its shielding effect is weaker. But geranium has a higher ionization energy than gallium because its effective nuclear charge is lower than in choice D. It states that Germanium has a lower ionization energy than silicon because its shielding effect is weaker. But Germanium has a higher ionization energy than gallium because its effective nuclear charge is higher. So with all of these choices outlined, we want to next refer to our periodic table to find the locations of Germanium and silicon. So in a second, I'm going to paste in our periodic table. So here we have our periodic table and we should notice that germanium and silicon are both in our same column or group number being four A on the periodic table where silicon we would observe being across period three and germanium we would or sorry gallium we would observe is across period four along with uranium. So let's note these observations down silicon and Germanium are both again group four A elements and across period. Or rather we'll say silicon is going to be in period three and Germanium is period four element. Whereas we find gallium in group three A and it's a cross period for along with Germanium. Let's recall that as we move down a group of the periodic table to higher subs shells or energy levels. So we'll represent energy levels by lower case N. This will correspond to a increase in our distance between the nucleus of an atom and valence electrons recall that valance electrons are outer electrons. So with these two notes in mind, with our increased distance between our nucleus of our atom and its outer electrons. And with an increase in subs shells between those two factors are nucleus and our outer electrons. Due to higher energy level, the inner electrons or the core electrons will have a higher shielding factor of the outer veins electrons and the positively charged nucleus in the center. So we would next have a correlation of an increased shielding effect. So an increase shielding between our core electrons from our valence electrons and the nucleus is going to allow for a weaker attraction a bounds electrons to the positive charged nucleus. And therefore, we would have an easier time removing an electron from the atom which will correspond to a decrease in our ionization energy of our atom as a whole. So the scenario where we would move down a group would be in our atoms in the same group being in four a on our periodic table where we have silicon and Germanium. And we would state that Germanium is in period four. And so therefore, it's lower than silicon. So therefore, Germanium has a lower ionization energy than silicon, which is, as we stated in period three. And this is because Germanium's shielding effect will use se for short, is going to be stronger based on that lower ionization energy. So it's easier to remove an electron from Germanium than it would be from silicon. So going to our answer choices, we're going to eliminate our choices which state that the shielding effect of Germanium is weaker. So we would rule out choices C and D. This leaves us with choices A and B and now we're going to consider our next two atoms which is between Germanium and gallium. Now, as we stated, Germanium and gallium are in different groups, but in the same period being period four So we're going to recall next and use a different color. We're going to recall that when we move from left to right across a period on the periodic table, which are called our rows of our periodic table. This corresponds to an increase in effective nuclear charge represented by now, we have an increase in effective nuclear charge because our orbital becomes more stable as it's being able to be filled with more electrons. And with a more stable orbital, we would have a increased amount of energy required to remove an electron. Since that preference would no longer be there as the orbital becomes more stable and an increase in energy required to remove an electron will correspond to an increased ionization energy of our atom. And as we stated, Germanium is more towards the right than gallium. So we'll say Germanium is a group for a element. So therefore, Germanium has a higher, has a higher effective nuclear charge then gallium, which as we stated is in group three A and it is because of this higher effective nuclear charge that Germanium has a higher ionization energy, then gallium, since again, there is more energy required to remove an electron from Germanium as it becomes more stable, being more further to the right in group four A than gallium is in group three A. So going back to our answer choices, we would eliminate between A and B choice B since it says that Germanium has a lower effective nuclear charge than gallium. We saw that it should be higher. So this leaves us with choice a which states that Germanium has a lower ionization energy than silicon because its shielding effect is stronger. We agree with that. But on the other hand, Germanium has a higher ionization energy than gallium because its effective nuclear charge is higher. And se A would be the best explanation for our phenomenon of why Germanium has a lower ionization energy than silicon but is found to have a higher ionization energy than gallium. So a is our final answer. I hope that this made sense and let us know if you have any questions.

Why is the ionization energy of Cl lower than F, but higher than that of S?

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Key Concepts

Ionization Energy

Atomic Structure and Electron Configuration

Electron Shielding