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Learning Objective Recognize the general periodic trends in ionization energy. Key Points The ionization energy is the energy required to remove an electron from its orbital around an atom to a point where it is no longer associated with that atom.
The ionization energy of the elements increases as one moves up a given group because the electrons are held in lower-energy orbitals, closer to the nucleus and therefore are more tightly bound harder to remove. Show Sources Boundless vets and curates high-quality, openly licensed content from around the Internet.
Licenses and Attributions. 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 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. 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. From this trend, Cesium is said to have the lowest ionization energy and Fluorine is said to have the highest ionization energy with the exception of Helium and Neon.
Each succeeding ionization energy is larger than the preceding energy. Electron orbitals are separated into various shells which have strong impacts on the ionization energies of the various electrons. For instance, let us look at aluminum. Aluminum is the first element of its period with electrons in the 3p shell. This makes the first ionization energy comparably low to the other elements in the same period, because it only has to get rid of one electron to make a stable 3s shell, the new valence electron shell.
However, once you've moved past the first ionization energy into the second ionization energy, there is a large jump in the amount of energy required to expel another electron. This is because you now are trying to take an electron from a fairly stable and full 3s electron shell.
Electron shells are also responsible for the shielding that was explained above. Both ionization energy and electron affinity have similar trend in the periodic table. For example, just as ionization energy increases along the periods, electron affinity also increases. Likewise, electron affinity decreases from top to bottom due to the same factor, i.
Halogens can capture an electron easily as compared to elements in the first and second group. This tendency to capture an electron in a gaseous state is termed as electronegativity.
This tendency also determines one of the chemical differences between Non metallic and metallic elements. Diagram 3: showing increasing trend of electron affinity from left to right 9.
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.
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