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Current time:0:00Total duration:7:32

Video transcript

like bonds are the bonds that hold together ionic compounds so basically it's what holds together cations and anions an example of a compound that's held together with ionic bonds is sodium chloride also known as table salt so here we have a close up picture of some really nice crystals of sodium chloride and this is something that you could try at home you can take some write any table salt that you might have dissolve it up in some water and then let that water slowly evaporate and if you're lucky you might get some beautiful symmetric crystals like these for me personally at least growing crystals that look beautiful is one of the most fun things about chemistry you can see when we when we look at the close-up shapes of these crystals that they have some very beautiful symmetry that symmetry tells us a little bit about the structure of these compounds on a molecular level if we zoom in on these crystals we can imagine actually we don't have to imagine you can look at these with different kinds of instruments like x-ray crystallography and you can look at the crystal lattice and get information about how the different ions are arranged in these solids so the way that the ions are arranged determines a lot of things about the properties of these compounds so these ionic bonds and how the ions are arranged tell us a lot about the solubility of the compound solubility and other properties like melting or boiling points and it even can be related back to things like how hard a particular ionic solid is so the ionic bonds here in the sodium chloride are the ones that hold together our sodium ions and our chloride ions so our sodium plus and our chloride minus and the strength of an ionic bond is related to the electrostatic force the electrostatic force between them and I'm going to abbreviate the electrostatic force as F subscript e so this is the force between two charged species and it's equal to some constant K times the two charges that are interacting divided by the distance between the two charges squared so here Q 1 and Q 2 are the charges and in the case of sodium chloride for example Q 1 and Q 2 would be Q 1 might be one plus from our sodium ion and Q 2 might be one minus from our chloride ion and we could also just switch those two we could say chloride is Q 1 and sodium is Q 2 and and that wouldn't change what we get from this equation and then R 2 here is the distance between the ions and we usually approximate it as saying it's the sum of the ionic radii for the two ions we're looking at so we can use Coulomb's law here to explain some properties that are related to the strengths of ionic bonds and so the example we're going to go through today is going to be that of melting point so we're going to look at some melting point trends and try to relate them to the different variables in Coulomb's law so the first thing we'll look at the first two compounds we'll compare our sodium fluoride and magnesium oxide sodium fluoride has a melting point of 993 degrees Celsius and magnesium oxide has a melting point of two thousand eight hundred and fifty-two degrees Celsius the other information we know about these two compounds if you look up the ionic radii it turns out that so D flouride the distance between the ions is about the same as magnesium oxide they're not exactly the same but they're pretty close so if we were to say that R is approximately the same for these two then we can explain the difference in melting points using the charges since melting point is a measure of basically how much energy do you need to add to these compounds to break apart your ion we would expect melting point to go up to increase as f:e increases as the force between the ions increases we would expect to have to add more energy to break those ions apart and we can see that in our first example magnesium oxide if we look at the charges on ions magnesium is two plus and oxide is two minus in sodium fluoride sodium is one plus and fluoride is one minus so we would expect assuming that R is about the same this q1 times q2 is four times bigger in magnesium oxide versus sodium fluoride so q1 and q2 product of q1 and q2 is higher for magnesium oxide and that's why we would expect the melting point to be higher we can also look at sodium chloride versus sodium fluoride and in this case let's look at well I don't know maybe this is kind of artificial the boiling point the melting point sorry the melting point of sodium chloride is 800 is 801 degrees Celsius and the melting point of sodium fluoride is like we said earlier 993 degrees Celsius and so this time the charges are the same on our ions our q1 and q2 is 1 plus for the sodium in both compounds and 1 minus for the chloride and the fluoride so q1 times q2 didn't change for these two compounds but since we change the anion from fluoride to chloride we increased our here and increasing R and the denominator makes the electrostatic force go down another way we could put it is that since R decreases as we go from sodium chloride to sodium fluoride the melting point goes up so in each of these pairs the compound that has a higher melting point is the one that also has the higher electrostatic forces and that's either because the charges are higher q1 and q2 are higher or because the distance between the ions went out so these are some examples for how we can relate the properties of ionic compounds to the electrostatic force using Coulomb's law between the cation and the anion
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