this is the section on molecular structure and polarity so the first we're going to take a look at the how to come up with the three dimensional shapes or structures of molecules and then we'll determine whether the molecule is polar or not so the system we're going to use is what's called an a b system okay a is the central atom and how many other atoms it's bonded to and eventually we'll have an e which will be the unshared pairs of electrons so you should have a chart in your slot nodes and the number of atoms the number of lone pairs that's not in your slot notes because that's just showing us how to come up with the a b system okay so we go to take a look at ab2 we go to take a look at beryllium chloride beryllium is found in column number two so it's got a plus two so it's going to have two bonds and because it's in group two it doesn't have any valence electrons left over to share so we end up with an ab2 and so we end up with an electron arrangement and a molecular geometry that are both the same and it's linear and this is 180 degrees we're not going to worry too much about that the degrees that's you would use for something else and so if we take a look we've got the ab2 now we're going to take a look at ab3 okay and so we're going to be taking a look at boron trifluoride okay boron's essential atom there were three fluorines around the outside and if we take a look at boron it only has three valence electrons so it's going to have zero bonded atoms around the outside and so its electron arrangement it's going to be trigonal planar and so is its molecular geometry because there's no unshared pairs of electrons if we take a look at the ab4 we're going to take a look at ch4 and if we take a look at carbon it's going to have four bonds or it's going to have four hydrogens bonded to it carbon only has four valence electrons four bonds so it's going to have zero lone pairs and that shape is going to be tetrahedral for both the electron arrangement and the molecular geometry ab5 is going to be trigonal bipyramidal and if you'll notice you've got x's in your slot notes because there's only five shapes that i want us to worry about the first three are linear trigonal planar and tetrahedral but i wanted you to see some of the other shapes that are out there and so the the ab6 which would have six bonds so we're going to put sulfur in the middle it's got six valence electrons if it shared all of them then it would have zero unshared pairs of electrons and it's going to be octahedral in shape so number question number seven explain the difference between electron pair geometry and molecular geometry okay so space must be provided for each pair of electrons whether they're in a bond or in a lone pair okay the electron pair geometry considers where the electrons are molecular structure considers the bonding pair geometry okay so it's looking where the atoms are bonded to it so up until this point all the electrons have been bonded so the electron arrangement and the molecular geometry have been the same but that's not always true so we're going to start out with looking at some of the ones we've already looked at and then look see what happens when we add some unshared pairs of electrons so our ab2e which means we've got two bonds and now we have an unshared pair of electrons if you'll notice our electron arrangement still trigonal planar but if we ignore the unshared pair of electrons what we end up with is we end up with a bent molecule notice that unshared pair of electrons that sticking out the top forces this molecule to go from linear to bent with our ab4 if you'll remember its electron arrangement would be tetrahedral because it would be four pairs of electrons but only three of there's only three of them that are being bonded and so we go to take a look at this shape if you notice an unshared pair of electrons which we're going to ignore that's why i'm kind of using the ball and stick models we're going to end up with what's called a trigonal pyramidal if we've got an atom that's got two bonds and two unshared pairs of electrons again because there's still four electron domains or four different groups of electrons two or two sets are in bonds and two sets are unshared pairs of electrons it's still going to be tetrahedral in shape but if we ignore those electrons what we find is we end up with a molecule that's bent and that's water and again there's only a certain number of shapes we're going to worry about i'm going to summarize them in just a little bit and so the next couple i just wanted to kind of go through and show you what they look like because they're actually exceptions to the octet rule and so if we do the ab5 system we can add in one unshared pair of electrons and so instead being trigonal bipyramidal we end up with the distorted tetrahedron it's also called a seesaw we can also now substitute in two unshared pairs of electrons and so we go from the trigonal bipyramidal that we started out with if we ignore those pairs of electrons in its shape what we end up with is we end up with a t-shaped if we now substitute in three unshared pairs of electrons we go from the trigonal bipyramidal original shape and now we're back to linear all those unshared pairs of electrons are just forcing the atoms to be in a straight line if we go on the octahedral shape the ab6 if we substitute in one unshared pair of electron and if we ignore the unshared pair of electrons sticking out the bottom of this we end up with a square pyramidal instead of the trigonal pyramidal and our last one ab4e2 so we got four bonds sticking out two unshared pairs of electrons so if we ignore those unchaired pairs of electrons those dots sticking out either end we end up with square planar instead of trigonal planar so if we use this system we can kind of identify the shapes of molecules so it says give the cbr4 so that's carbon carbon's in group number four and it's got four bonds for the bromines so it's an ab4e0 or an ab4 it's got zero unshared pairs so it's going to be tetrahedral in shape if we've got scl2 sulfurs in group number 16. so it's got six valence electrons it's got two bonds because it's got the chlorines and so we've got an ab2 and we've got 6 minus two is four four divided by two is two see how we're calculating the e2 and the ab2e2 is two lone pairs and it's going to be bent similar to our example of water we can do ions chlorines in the middle so it's got seven but we're going to add one i'm getting four bonds because remember oxygen's got a minus two so each oxygen is going to have two bonds so i'm going to have a two atoms bonded around the outside two unshared pairs of electrons and i'm going to end up with two lone pairs and it's going to be bent well we can have clo3 so if you'll notice now we've got six bonds three oxygens two bonds for each one so eight minus six is two two divided by two is one that's where i'm getting the e1 from is i have an ab3e1 and that's going to be the trigonal pyramidal if i have pcl3 phosphorus is in group 15 it's going to have three single bonded chlorines around the outside 5 minus 3 is 2 divided by 2 so we get e1 so i have ab3e1 and so that's going to be the trigonal pyramidal so if we take a look at question number 92 it's going to identify the electron pair geometry and the molecular structure i'm not going to ask you too much about the electron pair geometry i just want you to realize that there's a difference between the two so if there's no unshared pairs of electrons the molecular geometry and the electron geometry are the same carbon is in group number four it's got four fluorine so four minus four is zero so it's ab4 bf3 is trigonal planar becl2 is going to be linear because beryllium only has two valence electrons and there's two bonds connecting each of the chlorines so end up with a b2 so i end up with linear so if you'll notice i've given there's on your slide notes these are the five shapes that we're going to be interested in okay linear for co2 trigonal planar basically anything with boron in the middle of it boron trifluoride boron trichloride boron triiodide tetrahedral carbon with four atoms around the outside is going to be tetrahedral okay take a look at water it's bent and then the ammonia is trigonal pyramidal and these are just classic examples of the shapes so now what we're going to do is we're going to take a look at dipole moments and polar molecules dipole moments mean some moment in time the molecule is going to have an electron reach rich region and an electron poor region if the molecule at some point in time has a dipole moment or its die two different poles the molecule is said to be polar so this hf we've seen that in a previous slide if you'll notice the fluorine is electron rich and the hydrogen's electron poor because fluorine's got that really high electronegativity and so if you'll notice the hydrogen's positive and the fluorine is negative and if you'll notice we got the plus at the end of the arrow meaning it's positive and it's going to go from positive to negative where it's going to go from low amount of electrons to a high amount of electrons so if you'll notice we've got nh3 and nf3 they're both trigonal pyramidal but remember fluorine's got a higher electronegativity than the hydrogen so if you notice the dipole is kind of changing directions and the nitrogen is um more negative and so we went to 1.46 we can actually calculate it's not as easy as easy as it looks because we have to do the vectors in three dimensional space and time so calculating these out is not that simple but we get a general idea of the bottom of the molecule is going to be positive the top's going to be negative so it's going to have a dipole moment which means it's going to be polar if we take a look at our other molecule the fluorine is going to be more negative okay so if you'll notice there's going to be more negative for the fluorines how it's red towards the bottom of it and it's more red or more negative so because it's fighting the negative electrons we're going to end up with the smaller dipole moment the bigger the number the bigger the difference the greater the polarity so we're basically going to take three molecules we're going to try and predict whether each of the following molecule has a dipole moment or whether it's polar so one we could look up the dipole moments and we could calculate them out but i don't want to go that far in detail what i want to do is basically be able to look at a molecule and determine whether it's polar or nonpolar we're not going to determine how polar or how nonpolar we're just going to determine whether it could be polar or nonpolar so if we've got brcl we've got bromine and chlorine okay we know there's going to be a dipole because there's difference this bond is polar there's only two atoms so this one's relatively simple in that the chlorine is closer to fluorine so it's going to be more negative than the bromine and so it would have a dipole moment if we take a look at our bf3 notice our fluorines around the outside and so it's red all the way around the outside so it's going to be the same all the way around the outside so there's no dipole moment if we take a look at our lewis structure for methylene chloride notice this is ch2cl2 and we can't just look at this molecule and determine whether it's polar or not we need to think about its three-dimensional shape so if we put it in its tetrahedral shape what you'll notice that we end up with chlorines on one side and hydrogens on the other remember chlorine is closer to fluorine so there's going to be negative more negative than the hydrogen so notice the pluses are going the arrows and the hydrogens are going towards the carbon which means we're going to have a polar molecule that would be positive on the hydrogen or somewhat positive on the hydrogen and somewhat more negative on the chlorine side thus the ch2cl2 is a polar molecule so it says which one of the following molecules is non-polar well we're looking for something that would be the same all the way around the outside and so it's going to be the becl2 because it's going to be linear it says which of the following molecules has polar bonds but is a non-polar molecule which means it's going to have the differences in bonds but it's going to have the same all the way around the outside so that's going to be the bf3 it says which of the following species has the lot well remember fluorine is the most electronegative so i picked the one with the fluorine in it and it says which of the following molecules ions contain polar bonds they both contain polar bonds because they've got different atoms in it and they're both going to have a dipole moment which means they're both going to be polar