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  • - [Instructor] In this video we're going to look at trends

  • for the periodic table of elements

  • for dimensions like ionization energy,

  • atomic and ionic radii, electron affinity,

  • and electro negativity.

  • And to do so, we're going to start with

  • a very fundamental idea in chemistry or physics,

  • and that's Coulomb's Law.

  • And for our point of view, we can view Coulomb's Law

  • as saying that the magnitude of the force

  • between two charged particles is going to be

  • proportional, that just means proportional right there,

  • is going to be proportional to the charge

  • on the first particle times the charge

  • on the second particle, divided by the distance

  • between those two particles, squared.

  • When we're thinking about it in context

  • of the periodic table of elements and various atoms,

  • you can view q1 as the effective positive charge

  • from the protons in the nucleus of an atom.

  • You can view q2 as the charge of an electron.

  • Now any given electron is going to have

  • the same negative charge, but as we try to understand

  • trends in the period table of elements,

  • it's really the outer most shell electrons,

  • the valence electrons, that are most interesting.

  • Those are the ones that describe the reactivity.

  • And so when we think about the distance

  • between the two charges we're mainly going to be

  • thinking about the distance between the nucleus

  • and those outer most valence electrons.

  • Now we can view this effective charge,

  • I'll call it z-effective, as being equal to the difference

  • between the charge in the nucleus,

  • so you can just view this as the atomic number,

  • atomic number or the number of protons

  • that a given element or an atom of that element has,

  • and the difference between that and what is often known as

  • S, or how much shielding there is.

  • Now there is complicated models for that,

  • but for an introductory chemistry class,

  • this is often approximated by the number of core electrons.

  • Remember, we really want to think about

  • what's going on with the valence electrons.

  • And so if you imagine a nucleus here,

  • do that orange color, that has protons in it.

  • And so you have core electrons.

  • Let's say these are the core electrons in the first shell,

  • and then you have some core electrons in the second shell.

  • And let's say the valence electrons are in the third shell.

  • So let's say these are some valence electrons here,

  • they're blurred around, they're in these orbitals.

  • Those valence electrons, which have a negative charge,

  • are going to be attracted to the positive charge

  • of the nucleus but they're also going to be

  • repulsed by all these core electrons

  • that are in between them.

  • And so that's why an approximation

  • of the effective charge that these valence electrons

  • might experience is going to be the charge

  • of the nucleus minus, and this is an approximation,

  • the number of core electrons that you have.

  • So if we use that roughly as a way to think about

  • z-effective, what do you think are going to be

  • the trends in the periodic table of elements?

  • What would be the effective charge

  • for the Group I elements over here?

  • Well, Hydrogen has no core electrons

  • and it has an atomic number of 1.

  • So 1 minus 0 is going to have

  • an effective charge of roughly 1.

  • Lithium atomic number of 3, minus 2 core electrons

  • that are in 1-S, so once again you're going to have

  • 3 minus 2, effective charge of 1.

  • So roughly speaking, all of these Group I

  • elements have an effective charge of 1.

  • What if you were to go to the halogens?

  • What's the effective charge there?

  • Well if you look at Flourine, atomic number of 9,

  • has 2 core electrons in the first shell,

  • so has an effective charge of 7.

  • Chlorine actually has an effective charge of 7

  • for the same reason.

  • Atomic number of 17, but 10 core electrons.

  • If you go even further to the right,

  • to the noble gases, you see that Helium

  • is going to have an effective charge of 2,

  • atomic number of 2 minus 0 core electrons.

  • But then when you get to Neon,

  • you have an atomic number of 10,

  • and then minus only 2 core electrons.

  • And you'll see as you go down these noble gases,

  • other than Helium, they have an effective charge of 8.

  • And so the general trend is, your effective charge is low

  • at the left, effective charge low for Group I,

  • and then when you go to the right of the periodic table,

  • you have a z-effective, is going to be high.

  • So within a given period, or within a given row

  • in the periodic table of elements,

  • your outer electrons, your valence electrons,

  • are in the same shell.

  • But the effective charge is increasing

  • as you go from left to right.

  • So this q1 right over here is going to be increasing.

  • So what is that going to do to the radius of the atom?

  • Well, Coulomb's Law will say that the magnitude

  • of the attractive force between those opposite charges

  • is going to be stronger.

  • And so even though you're adding electrons

  • as you go from left to right within a row,

  • within a period, the atoms in general

  • are actually going to get smaller.

  • Let me write it this way.

  • So as you go from left to right, generally speaking,

  • radius decreases.

  • Now what's the trend within a column?

  • Well one way to think about it is,

  • as you go down a column, as you go down a Group,

  • you're filling shells that are further out.

  • And so you'd expect radius to increase

  • as you go down a column, or down a Group.

  • Or you could say radius decreases as you go up a group.

  • So radius decreases.

  • So overall what's the trend

  • in the periodic table of elements?

  • Well radius is going to decrease as you go

  • up and to the right.

  • And so you could draw an arrow something like this.

  • And it is indeed the case that by most measures,

  • Helium is considered to be the smallest atom,

  • a neutral Helium atom.

  • And Francium is considered to be the largest atom.

  • So could we use this to think about

  • other trends in the periodic table of elements?

  • What about, for example, ionization energy?

  • Just as a reminder, the first ionization energy

  • is the minimum energy required

  • to remove that first electron

  • from a neutral version of that element.

  • And since it's the minimum energy,

  • it's going to be one of those outer most electrons.

  • It's going to be one of the valence electrons.

  • And so what's going to drive that?

  • Well you can imagine the ionization energy

  • is going to be high in cases where

  • the Coulomb forces are high.

  • And what are the situations where

  • the Coulomb forces are high?

  • Well this is going to be a situation

  • where you have a high effective charge

  • and where you have a low radius.

  • Low radius makes the Coulomb forces high.

  • And effective charge makes the Coulomb forces high.

  • So where is that true?

  • So you have the lowest radii at the top right

  • and you have the highest effective charge at the right.

  • So you would expect the highest ionization energies

  • to occur in the top right.

  • So high ionization energy.

  • And that actually makes intuitive sense.

  • These noble gases are very stable.

  • They don't want to release an electron.

  • So it's going to take a lot of energy

  • to take one of those electrons away.

  • Fluorine or Chlorine, they're so close

  • to completing a shell, the last thing they want to do

  • is lose an electron.

  • So once again, it takes a lot of energy

  • to take that first electron away.

  • On the other hand, if you go to something

  • like Francium, it has one valence electron.

  • And that valence electron is pretty far from the nucleus.

  • And there's a low effective charge