Periodic Trends - Ionization Energy
Topics Covered: Periodic Trends, Coulomb’s law, effective nuclear charge, shielding, ionization energy
Have you ever looked at the periodic table and wondered, why are the elements organized in that way? What’s so special about the periodic table? Well, in this post and the series of periodic trends that will follow, we’ll cover the concept of the periodic trends: the property that changes from top to bottom, and left to right of the periodic table. And today, we will begin the series by talking about the ionization energy.
Before we begin our discussion, let’s start off by defining some terms that will be useful later.
Most of the trends on the periodic table can be explained in four concepts: Coulomb’s law, effective nuclear charge, shielding effect, and electron-electron repulsion. Today, we’ll only cover Coulomb's law, effective nuclear charge, and shielding effect.
Coulomb’s law is the law that states the following: F = kq1q2 / r2, where F = force of attraction or repulsion, k = Coulomb’s constant, q1 & q2 = charges, r = distance between the two charges. Looks complicated? Don’t worry, there will rarely be any case in which you have to plug in numbers and figure out the actual force of attraction or repulsion. All you need to know is the relationship between charges and force, and distance and force. See that q1 and q2 are at the numerator? That means as the charges increase, the force of attraction or repulsion will also increase. How about distance? Well, r is at the denominator of the fraction, so as the distance increases, the force will decrease!
Effective nuclear charge is the net positive charge experienced by the electron. In other words, it is the amount of attraction that outermost electrons experience by the nucleus. Written as Zeff, effective nuclear charge is calculated as Z - S, where Z is the number of protons and S is the number of shielding electrons (electrons that are not valence electrons)*.
The shielding effect is when the inner electrons “shield” the nucleus and prevent the valence electrons from the full attraction of the nucleus. Think of the shielding electrons as a barrier. The positive protons in the nucleus try to attract the valence electrons that have negative charge, but the barrier of shielding electrons blocks that attraction and reduces the force of attraction between the nucleus and valence electrons.
Ok, now let’s dive into the concept of ionization energy. Ionization energy is the energy required to remove an electron from an atom. So what’s the trend like? Let’s look at the periodic trend first. Generally**, as you go from left to right of the periodic table, the ionization energy increases. This can be explained by effective nuclear charge. As you go from left to right, the number of protons in each atom increases (while the number of inner electrons stay the same), so the effective nuclear charge Z -S would increase. This means the nucleus will be attracting the valence electrons harder, and it would require greater energy to remove an electron from an atom!
Let’s look at the group trend. As you go from top to bottom of the periodic table, the ionization energy will decrease. This is due to the shielding effect. As you go from top to bottom, the sizes of atoms increase (more energy levels being added), and the size of electron clouds and the number of shielding electrons would increase as well. Therefore, due to shielding effect, more core electrons would block the attraction between nucleus and valence electrons, so the valence electrons would be held more loosely to the atom. Therefore, it would require less energy to remove an electron from an atom!
Finally, let’s explain these trends with Coulomb’s law as well. Going from left to right of the periodic table, the number of protons (charges) increase, so the force of attraction will also increase. Going from top to bottom of the periodic table, the distance between the nucleus and valence electrons increases (as atom size gets bigger), so the force of attraction will decrease. Now you can explain the periodic trend of ionization energy!
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* In this post, I will be talking about inner electrons, core electrons, and shielding electrons, but they are all referring to the same thing —electrons that are not valence electrons!
** I say generally because this trend isn’t ALWAYS true. For example, although oxygen is on the right side of nitrogen, it has less ionization energy than nitrogen because oxygen’s repulsion of the paired electrons from its electron configuration makes it easier to remove an electron from oxygen than from nitrogen. As always, chemistry is full of EXCEPTIONS, which is what makes chemistry more interesting! :)
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