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Ionization energy increases across a period because there is a higher effective nuclear charge, so the electrons are held more tightly, and thus requires more energy to remove an electron. Ionization energy also decreases down a group, since the outer electrons are farther from the nucleus and are held less tightly, requiring less energy to remove an electron.
The other responses already did a great job describing the trend, but if you are comparing elements that are close together on the period table or the exceptions, creating an aufbau diagram will be helpful. Since ionization energy measures the amount of energy required to remove an electron, having a half filled or filled orbital is favorable, so removing an e- will require more energy. This is seen between the ionization energies for oxygen compared to nitrogen and beryllium compared to boron. The reasoning behind these exceptions are clearly visible with a diagram and you can see the favorable vs less favorable electrons interactions.
As said before, the general trend is that ionization energy increases from down to up and left to right on the periodic table. However, keep in mind that nitrogen has a higher ionization energy then oxygen and beryllium has a higher ionization energy then boron. Hope this helps!
The reason ionization energy increases from down to up and left to right on the periodic table is due to energy shells and shielding of electrons. To be more specific, as you go right across a period, the shell in itself is not increasing. So the electrons do not get farther away from the positively charged nucleus. In other words, you can think of the electrons as staying a constant distance away from the nucleus, even though I'm pretty sure it isn't scientifically correct. However, the protons the atom are increasing, meaning that the pull from the nucleus is going to be greater, thus a smaller ionic radius, thus a larger ionization energy. It takes more energy to pull the electron away from the attraction from the positively charged nucleus. The opposite relates to going down a group. As you go down a group, the energy levels are increasing, thus the states the electrons could be in are further away from the nucleus. As a result, the ionization energy is way smaller, because the valence electrons don't need a lot of energy to escape the pull of the positively charged nucleus. Really just typed an essay here
Ionization increases across a period since the atomic radius is smaller and the electrons are more tightly held thus requiring more energy to remove the electron. It decreases down a group since the atomic radius is increasing and electrons aren't as tightly held and are shielded from the positive charge of the nucleus making it easier to remove an electron.
I find it helpful to remember that F has the highest ionization energy. From there, I know that almost every trend moves diagonally. So if F has the highest ionization energy, the trend Increases moves up and to the right.
another good way to remember is to contrast this with electronegativity. Electronegativity means that an atom has a strong ability to attract electrons, meaning that it would take a lot more energy to remove one if we wished to. In other words, a high electronegativity results in a low ionization energy, and vice versa. Try to remember just one way, since it may be confusing to try to remember all of the trends. Hope this helps!
Ionization energy increases as you go across a period and decreases as you go down a group. As you go across the period the number of protons in each element increases, which draws the electrons in more tightly and makes them more difficult to remove. As you go down a group, the valence electrons are occupying shells that are further away from the nucleus; the attraction of the nucleus is weaker and makes the electrons easier to remove.
Ionization energy increases going right on the periodic table because the number of protons in the nucleus increases, the ionization energy or attraction will increase. The nucleus brings in the electrons closer together because of the increase in attraction. Ionization energy will then decrease going down the periodic table because of the increased distance between the valence electrons and the nucleus, and the nucleus will not be able to pull those electrons closer.
I like to think about periodic trends in the way that makes sense instead of memorizing how they move, there are quite a few trends and it can be easy to get them muddled if you do not think about how it may work. Ionization energy is the amount of energy required to remove an electron from a gaseous element. If you think about it, it is much more favorable to remove an electron from a group 1 element because to get to a noble gas state an electron has to be removed, however in elements like fluorine, it is very unfavorable to remove an electron as they need to gain electrons to have the noble gas configuration, therefore it will take more energy to remove an electron. So, the ionization energy increases across a period. Next, if you consider effective nuclear charge and its hold on the last electron, it makes sense why ionization energy would decrease down a period. The electron being removed from cesium has a lot less pull from the nucleus than an electron removed from lithium. Thus, lithium requires more energy to remove its outermost electron than cesium.
The ionization energy increases going to the right because the atoms become more electronegative and want to hold onto their electrons more tightly. It decreases going down a group because the electrons in the outer shell are not held onto as tightly and can be removed more easily.
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