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  • - [Instructor] You may already be familiar

  • with Coulomb's law,

  • which is really the most important

  • or underlying law behind all of what we know

  • about electrostatics

  • and how things with charge attract or repulse each other,

  • but a simplified version of Coulomb's law

  • is just that the force between charged particles,

  • the magnitude of the force

  • is going to be proportional to the product of the charges,

  • so q one would be the charge

  • of one of the charged particles.

  • Maybe this is an ion.

  • Q two would the charge of the other particle.

  • Maybe that's an ion,

  • divided by r squared.

  • And if we're talking about ions,

  • r is going to be the distance between their nuclei,

  • and if the charges are different,

  • it's going to be force of attraction.

  • If the charges are the same,

  • it's going to be a force of repulsion.

  • And we can use Coulomb's law to think about ionic compounds.

  • So let's go with maybe the most common ionic compound

  • in our daily life, and that is table salt.

  • Table salt is sodium chloride,

  • so sodium chloride.

  • We have talked about this in other videos.

  • It is made up of positively-charged sodium cations,

  • so you have an Na plus,

  • so sodium is a group one element.

  • It's very easy to nab an electron off of it

  • and then it has a positive charge,

  • and it's made up of a chloride anion,

  • so Cl minus.

  • Chloride is a group seven element.

  • It really wants to get that extra electron

  • to have eight valence electrons in its outermost shell,

  • and so it's very likely to grab an electron maybe

  • from a sodium,

  • and so these two characters are going to be attracted

  • to each other.

  • Notice, they have opposite charges.

  • And when you have a bunch of sodium and chloride together,

  • you'll have a structure that looks something like this.

  • And in chemistry, we call this a lattice.

  • Now in everyday language,

  • you might associate things like lattices

  • with kind of a crossing pattern like that,

  • and in chemistry, when we're talking about a lattice,

  • we're talking about a three-dimensional structure

  • of atoms or three-dimensional structure of ions

  • that have a repeating pattern to them,

  • and you can see that here,

  • and in future videos, we'll go into more detail

  • onto lattice structures,

  • but you can see in this picture,

  • the purples are the sodium cations

  • and the greens are the chloride anions.

  • And the reason why the sodium cations are so small,

  • you can see that if you look at the periodic table

  • of elements here.

  • We have said that as you go to the right,

  • your radius decreases,

  • but what's happening is when sodium loses

  • that outermost electron,

  • then its electrons have a noble gas configuration of neon.

  • So it really loses that third shell, it gets smaller,

  • and not only does it lose that third shell,

  • but it has 11 protons,

  • so it's going to have a very strong pull

  • on those electrons in that second shell.

  • And similarly, chloride is going to gain an electron

  • so it's going to have a noble gas configuration of argon.

  • So it is going to be bigger.

  • Now when we talked about covalent bonds,

  • we talked about the bond energy,

  • the energy needed to pull apart the atoms

  • that were forming the covalent bonds.

  • There's a similar notion for ionic bonds like this

  • and that is lattice energy,

  • and that is energy necessary

  • to pull the ions apart

  • so that they are infinitely far apart from each other,

  • and lattice energy is usually measured

  • in kilojoules per mole,

  • which is also what we measure bond energy in

  • because they're really the same notion,

  • except lattice energy,

  • you're breaking up a lattice of ions,

  • while in bond energy,

  • you're normally talking about covalent bonds.

  • Now I want you to think about something.

  • What's going to have a higher lattice energy?

  • Would it be sodium chloride,

  • or let's pick something else.

  • Let's say we had rubidium.

  • Rubidium chloride,

  • which is going to have a higher lattice energy?

  • What's going to take more energy

  • to pull the ions apart?

  • And I'll give you a hint

  • with this periodic table of elements.

  • All right, well, rubidium chloride,

  • that's made up, instead of a sodium cation,

  • that's made up of a rubidium cation,

  • so you have Rb plus, and of course,

  • you have the chloride anion, Cl minus,

  • and so what's the difference here?

  • The anion is both, is chloride in both cases,

  • but when you look at rubidium versus sodium,

  • rubidium, when it loses an electron,

  • it's going to have a noble gas structure,

  • electron structure of krypton,

  • while sodium, once it loses an electron,

  • it's, its electron,

  • its electron configuration is going to look like neon.

  • So the sodium cation is smaller,

  • and what does that tell us?

  • Well, if this one right over here,

  • let me circle it like this.

  • If this is smaller,

  • and we have similar charges on top,

  • you have a plus one and a negative one on top,

  • that's the charges between the two ions,

  • but now you have a smaller radius between the nuclei

  • because sodium is smaller than rubidium.

  • While the radius goes down,

  • the force goes up,

  • so you're going to have stronger Coulomb forces

  • in a lattice of sodium chloride

  • than in a lattice of rubidium chloride.

  • Because the force of attraction is stronger,

  • it's going to take more energy to pull it apart.

  • So because of that,

  • you're going to have a higher,

  • higher lattice energy.

  • Lattice energy for sodium chloride than rubidium chloride.

  • Let's think about another ionic compound.

  • Let's say we were to think about magnesium fluoride, F two,

  • and this is made up of a magnesium cation

  • that has a positive two charge,

  • so two plus, in a lattice with a bunch of fluoride anions,

  • so with a bunch of fluoride anions.

  • So how would the lattice energy

  • of magnesium fluoride compare

  • to what we just saw up here?

  • So magnesium has a larger charge

  • than these cations up here,

  • so if you viewed the charge of magnesium as q one,

  • you're going to have something larger up there

  • and that fluoride is a smaller anion than chloride.

  • We can see that if we look at the periodic table

  • of elements again.

  • Florine is smaller than chlorine,

  • and so even if you added an electron to both of them,

  • fluoride is still going to be smaller,

  • and magnesium, when you take two electrons off of it,

  • it's going to have the noble gas configure,

  • electron configuration of neon,

  • but it's going to pull even more on those,

  • that, those second shell electrons

  • because it has 12 protons versus sodium only has 11.

  • So what we see here is not only does magnesium have a larger

  • positive charge than the sodium cation does,

  • but it's going to be smaller.

  • And the fluoride has a comparable charge to the chloride,

  • but it too is going to be smaller.

  • So we have a larger charge on top,

  • at least for the magnesium,

  • and you have smaller radii for the bottom,

  • so in magnesium fluoride,

  • the Coulomb forces between the ions and the lattice

  • are even stronger,

  • and so the lattice energy,