Hi. It's Mr. Andersen and this is chemistry essentials video 19 on covalent bonding. When you look at these three crystals over here, they look somewhat similar.
So what here. But the bonds that hold them together are going to be really different. We've got metallic bonding up here.
Ionic bonding here. And then we have covalent bonding here. And so what's going on with the electrons is the best way to tell each of these apart.
And when we start with covalent bonding you really have to start by talking about electronegativity. Which is how much an atom wants an electron. And it's going to increase as we go up and across on the periodic table.
And the reason why, it's based on the shell model and Coulomb's law. And so as we decrease the size of an atom, as they get smaller, it's going to be closer to the protons that are found in the nucleus. And then as we move over to the right those positive values are going to increase. And we're also going to increase that attractive value. And so based on those electronegativities we can look at the atoms inside a molecule and we can figure out is it a covalent bond?
And if so, what type of covalent bond is it? Now if those electronegativities are exactly the same then we call that a nonpolar. covalent bond.
An example would be oxygen that you're breathing right now. It's two oxygen atoms. They have the same electronegativity.
So they're going to share the electrons equally. Also you should know this. If it's a carbon and a hydrogen that are attached together, carbon is really good at sharing its electrons.
And so we generally call that a nonpolar covalent bond. We can look at these bonds using a graph where we graph the potential energy on one side and then the distance between the atoms on the other side. And so what becomes apparent is going to be the bond length, how far those atoms are apart and then the bond energy.
How much energy would we have to put in to break those atoms apart? Now if the electronegativity is different then we call that a polar covalent bond. And so when we do that we have something called a dipole where there's a partial positive and a partial negative charge.
Whichever one has the highest electronegativity is going to have that negative charge. But But even though we may have differences in charge, we still have conservation of charge over that whole molecule itself. Also as we move from covalent bond to more and more electronegativity we start to move into what are called ionic bonds.
And it's a gradient. It's a continuum. And so we can measure these to tell the difference between them.
And so a covalent bond remember is going to be sharing of the electrons. So a great example would be hydrogen gas. And Hydrogen has one electron in its outer shell. It would like to have two. And so it can share that one electron with another hydrogen atom.
And now we have the sharing of these two electrons. And so we have a covalent bond between the two. Carbon for example has four valence electrons. According to the octet rule it would like to have eight.
And so what it can do is it can share each of those four electrons with another hydrogen atom. And now we have methane which is going to be. a non-polar covalent bond. And so here's that electronegativity.
Remember as we go up on the periodic table it's going to increase. And then as we go across it's going to increase as well. Now why is it increasing when we go up? It's because the atoms are becoming smaller. And as the atoms become smaller, those electrons are closer to the nucleus and so there's going to be a greater charge.
And as we move across you can see the electronegativity is increasing as well because we're increasing the amount of positive charges on the inside. And so we're increasing that pole. And so you can see that fluorine is going to have the highest electronegativity on the periodic table.
That means it wants electrons more than any other atom. And so if we have two atoms that are sharing electrons and the electronegativity differences are zero, we call that nonpolar. So oxygen for example. If we look at oxygen, its electronegativity is 3.44. And so we're going to have a polar polar.
And so we're So if we compare it to itself it's going to be no difference between the two. So that would be a non-polar bond. Same thing with nitrogen. Remember if we're ever connecting carbon and hydrogen, they're going to share the electrons equally. And so we'd also call that a non-polar covalent bond like in methane.
Now we can study these using a graph like this where we put potential energy on one side and then we put inner nuclear distance between the two. So let's start with hydrogen or H2. If we have two hydrogen atoms together, As we move them really close together they're going to want to push on each other.
In other words there's going to be repulsion. As we move them farther apart eventually they're going to get into an area where there's this sweet spot between attraction. And then as we move them really far apart they're going to be so far apart that Coulomb's Law says there's not going to be a great enough force to pull them together.
And so if we're below this line it's attraction. But what we'll find is when we have the perfect distance apart. We're going to reach what is the bond energy and the bond length of those two atoms.
In other words bond length is going to be how far generally hydrogen atoms are going to stay apart. And bond energy is how much energy would we have to put in to actually break that hydrogen apart. And so this graph is incredibly important.
When we're looking at nonpolar covalent bonds what we're going to find is that their electronegativities are different. So this right here would be hydrochloric acid. If we look at chlorine 3.16 and hydrogen 2.20 There's a difference between those electronegativities.
And so the chlorine is actually going to be a negative charge and the hydrogen is going to have a positive partial charge. Or if we look at water, oxygen at 3.44, hydrogen at 2.20. And so the oxygen is going to pull those electrons closer to it. So you're going to have a partial negative charge here and then a positive charge out here.
And that really allows us to have water attracting the hydrochloric acid. There's going to be a dipolar charge and the connection between those two would intermolecular at this point. And so how do you figure out what type of bond it is? Well there's one quick way to do it. You could just look at their electronegativity differences on that chart.
And if it's less than 0.5 it's generally going to be non-polar. If it's between 0.5 and 1.7 it's going to be polar covalent. And if it's greater than that, generally it's going to be ionic. And we'll talk about that in the next video. That's not a great way to do it.
Better ways are to look at what's being connected. Generally in covalent bonds it's going to be two nonmetals. In ionic it's going to be metal and nonmetal.
And then more important than that we should look at the properties. And so if we're looking at covalent bonds generally we're going to have gases, liquids, solids. Low melting and boiling point and they're going to be poor conductors.
But as we move to ionic it's going to be a crystalline solid. High melting boiling points and we're going to be good conductors especially when we dissolve them in water. And so did you learn the following to rank polarity based on location of atoms in the periodic table? Remember as we go up you're going to see that the And to the right we're going to increase electrode negativity based on Coulomb's law and the shell model.
And I hope that was helpful.