hello everyone it's mrs hansen once again picking up with our lessons in our chapter one of organic chemistry section five called polar covalent bonds we understand bonds a shared pair of electrons can be one of several categories it can be a non-polar covalent bond where the electronegativity difference is less than 0.5 when you see me write e n just know that that stands for electronegativity and again review from your probably your first semester gen chem those examples if i look at carbon's electronegativity value of 2.5 and i subtract the exact same number we're going to have an electronegativity difference of zero any element bonded to itself would be a nonpolar covalent bond and even carbon which is 2.5 with a hydrogen which is assigned a 2.1 gives us an electronegativity difference of 0.4 which is a very negligible difference they are considered to be non-polar covalent bonds if we have an unequal pair that unequal electron distribution we have a polar covalent bond the electronegativity difference would be between 0.5 and 1.7 carbon with its 2.5 electronegativity value compared to oxygen which is 3.5 gives an electronegativity value of 1.0 which lies in the polar covalent region notice that we use special designations to indicate the induction of charge carbon to oxygen we use this tail at the end of an arrow and we point to the more pa the more electronegative element carbon is less electronegative resulting in a little positive charge and oxygen which is more electronegative more electron rich gives it a partial negative charge so here we can see electron deficient gets that tail portion of the arrow or the positive symbol an electron rich gets that negative region and we can represent polar covalent bonds with those types of symbols and if the electronegativity difference is so large greater than 1.7 it's not covalent it's not sharing at all but it's more of an electron transfer and that indicates that we have ions notice that the 0.9 assigned from sodium when compared to oxygen which is 3.5 and a plus an o h minus would be a sodium ion and a hydroxide ion giving an ionic bond so based on electronegativities we're able to figure out what type of electron shearing is occurring sometimes we have an electronegativity difference that is right on the edge of these arbitrary signs of being polar covalent or ionic and as an example we see a carbon to lithium bond when i take a carbon which if you remember is assigned 2.5 and lithium who is assigned a 1.0 just to verify lithium is 1.0 we can see that that electronegativity difference is 1.5 which is very close to the 1.7 that is arbitrarily selected as being ionic and so what is correct do we show it as a shared pair of electrons as a covalent bond or do we show it as ionic with the carbon which is the more electronegative having a negative charge remember that's its formal charge and lithium with its plus one and the answer is yes they're both acceptable to write either way polar covalent or ionic and just keep in mind that this line is pulling electronegativity values is just one of many methods for calculating these differences so if it's close you're not going to be marked incorrect because both are acceptable if you write one or the other both are correct more of this skill in this level is to identify polar covalent bonds and result from inductive effects placing those partial charges and so the partial charges just remind you could be the arrow which the positive tail is going to represent the more positive region and then the more electronegative element is being pointed at or you can use these little greek symbols for partial positive or partial negative both of those symbols are used to show polarity inside of a bond so let's find all the polar covalent bonds in this structure called methanol and remember we had just said a carbon to carbon and a carbon to hydrogen you should just flat out memorize these are non-polar so right off the bat i can eliminate those three bonds as not being polar here i have a bond between carbon to oxygen and that clearly has an electronegativity difference that falls in the polar range and who is the more electronegative element ask yourself who's closer to fluorine and that would be oxygen and so the bond polarity is pulling the electron density towards the oxygen leaving the carbon electron deficient with a partial positive charge and the oxygen as electron rich giving it a partial negative charge and there's also a polar bond between the hydrogen and the oxygen the oxygen again more electronegative is pulling the electron density towards itself leaving the partial positive charge on the hydrogen as well so that induction the withdrawal of electrons towards the more electronegative element is shown either with these partial positive or partial negative signs or with the polarity bond arrow should we try another that's our answer from before we had said methanol could be used with an arrow or with a partial positive partial negative and so we had just drawn those but let's try some more from homework so i just selected a few from your skill builders and right off the bat just notice that the carbon to hydrogen are nonpolar so take those right off the plate we're not going to have to indicate any type of polarity there but i do see a carbon to oxygen bond in both of these directions noticing that the carbon is electron deficient the oxygen would be electron rich so i can see the partial negative on oxygen in each of these partial positives on the carbons this is a mirror image we have a carbon to oxygen bond here making this bond electron deficient at the carbon an electron rich at the oxygen and now we have all of those symbols placed in there to represent the partial positives and partial negatives in those polar bonds so i see four polar bonds all of them being the same carbon to oxygen letter b c h are nonpolar so they're uninteresting to us the carbon to fluorine is a very polar bond making the fluorine who is the most electronegative element very electron rich and the partial positive left on carbon same idea here carbon having its electron density being pulled away from it towards the chlorine making the chlorine electron rich and the carbon electron deficient and in letter c i also recognize that there are two polar covalent bonds the first between magnesium and bromine bromine of course is more electronegative than magnesium making bromine very electron rich and magnesium very electron deficient positively charged carbon is more electronegative than the magnesium as well making carbon here electron rich in comparison to magnesium we have identified two polar covalent bonds in structure c we've identified two polar covalent bonds in structure b and all four polar covalent bonds in structure a our next section reviews atomic orbitals those s p d and f shaped orbitals from quantum mechanics this is the section where you learned how to write electron configurations 1s2 2s2 and so forth and so just a little quick review of quantum mechanics back in the 1920s was the established theory that explains the wave properties of electrons the solution to these wave equations gave us wave functions and these functions which were psi squared gives us these images of atomic orbitals we know that an s-shaped orbital is a sphere and a p shaped orbital looks like an hourglass if the hourglass these p shaped orbitals are aligning itself on the y axis of an x y z three dimensional we called it p sub y if they're aligning itself on the x axis we called it p sub x and p sub z is our third spatial orientation for p orbitals notice that these p sub x p sub y and p sub z are all of equal energy and we need to remember that vocabulary word called degenerate means they're of equal energy an s is of lower energy than a p which is of higher energy but those p orbitals are of equal energy to one another so those p orbitals are called degenerate when we're filling atomic orbitals with their electrons there were three rules or three guiding principles that allowed us to place electrons the first of those three is called the aufbau principle simply stated the aufbau principle says the lowest energy orbital is always filled first keeping in mind that that means that the 1s orbital is always filled with at most two electrons the 2s orbital fills next with at most two electrons the 2p orbital fills px pypz for a total of 6 electrons and so forth the polyexclusion principle reminds us that each orbital can accommodate the maximum of two electrons but they must have opposite spin and that means as we were writing our orbital diagrams we would place one electron first and then partner it with an opposite spin we would place one electron first and then partner it with an opposite spin hund's rule when dealing with degenerate orbitals such as the two p sub x sub y sub z one electron is placed into each orbital first before you partner up so i had just modeled that if i'm filling the 2s orbital it fills completely then i start placing in the p's one in each first and then assign partners with opposite spin most students do very well in that section of gen chem it'll come back to you with some quick practice when we write electron configurations let's review so one of the tricks that i like to use is i don't know if it's a trick but it's just a visual to help me remember i like to write the orbitals and i'm just going to kind of create a column that says in the first column i'm going to have all the s's and i remember that s is fill with at most two and at the first energy level only has one sublevel so i'm done with that the second energy has two sublevels so i'm going to just continue writing 2 p 3 p 4 p in a column and i know that at most the p orbitals hold six electrons and in a third energy level we have three sublevels and the d's can hold up to 10 and the fourth energy whoops the fourth energy level can have four sublevels and at most 14. so that's kind of just a rough idea and just kind of saying i know that i'm going to fill 1s2 i hit the end of the row so i come back to the top 2s2 come to the end of the row so i come back to the top 3p 2p3s 3p4s 3d4 p5s and so forth and so that little trick helps to remember the order of which we fill and again that should come back to you rather quickly now on the periodic table you find that carbon is atomic number six so we know to write for six electrons the first of those six electrons are written in the one s2 2s2 that gives us four of the six and then we have two p2 gives us a total of six electrons so the orbital diagram 1s2 2s2 2p2 where i have an empty orbital and two unpaired electrons for oxygen oxygen is number eight on the periodic table so we have eight electrons to write for 1s2 2s2 2p4 gives us a total of eight electrons you can see in the outermost energy level there's two electrons in the s four more in the p giving us one that is partnered and two that are not in the p sublevel and boron which is element number five 1s2 2s2 2p1 giving us a total of three valence electrons two of them in the s subshell one more in the p so a total of three valence electrons about if we see positive or negative charges remember that a negative charge means that it has an extra electron if it has a positive charge you have to subtract an electron because it lost that electron so a negative charge one extra electron well up above here carbon is an atom started as 1s2 2s2 2p2 but if we added an electron instead of having 6 we now have 7 electrons meaning that we'd have 2p3 as its total number of electrons so carbon with an extra electron has an electron configuration with one more electron than its neutral atom same idea with carbon as a positive it means it lost one electron so if it had six and it lost one it means now that we're going to have five electrons so lost one gives us 5 gained 1 gives us 7 up above nitrogen on the periodic table is number 7 so 1s2 2s2 2p3 would give us the seven electrons but now that it's positive it lost one electron so we only write for six electrons 1s2 2s2 2p6 is the nitrogen plus one by the way that vocabulary word is called isoelectronic with carbon in other words a carbon atom and a nitrogen positive both have the same number of electrons and that would be six electrons and this is just a quick visual of what we've been practicing hydrogen with its first electron 1s1 helium with its two valence electrons 1s2 lithium has three valence electrons the first two in the 1s sublevel one more in the 2s sublevel beryllium which is number 4 1s2 2s2 boron ago and so forth so writing those electron configurations is review from this section in the next section we talk about covalent bonding and the overlap of atomic orbitals and this is known as the valence bond theory when electrons are treated as waves and remember we talked about one hydrogen atom kind of overlapping another hydrogen atom and that's what they're kind of showing here as they begin coming closer and closer together the internuclear distance starts bringing them closer together and those waves pass over each other with a constructive constructive interference we have a bond overlap that forms for a covalent bond if the bonds come together and they are in a destructive pattern where they cancel each other out we get a node and that's known as destructive interference and so literally we just want to kind of envision that the covalent bond is bringing this would represent an s-shaped orbital overlapping with an s-shaped orbital forming this covalent bond and this could represent for instance one hydrogen to the next an overlapping orbital the valence bond theory when those bonds overlap and those the major part of the density is located on the bond axes this is what we called a sigma bond all single bonds are sigma bonds so for instance if an s overlaps an s orbital it's lying on the same axis and s s overlap as a sigma bond i could have an s-shaped cloud overlap a p-shaped cloud head-to-head and it would still have of course i should have drawn those actually overlapping that would have been better but it still has all of the electron density on the x axes so an sp overlap could also be a sigma bond and if i kind of draw that a little nicer here are some examples of sigma bonds as i tried to draw an ss two spherical shapes overlapping in the same electron density as the x axes here's an sp here would be a pp and sp hybridization notice that even 2p clouds could also overlap head to head along the same internet axis so here we all have examples of sigma bonds which means that the overlapping orbitals lie within the same x axes and when i say that it's the same axes as the nuclei right that's where the nuclei resides is in the center of the atom so it's along the same electron density is where you find the nuclei of the atom a pi bond if you remember occurs when orbitals overlap that are either above or below the axes of the nuclei so if it's on the same axis we called it a sigma but if it's above the axes or below the axes where the pi cloud forms from the p orbitals above and below we call that a pi bond remember the sigma bond are all single bonds as well as the first in a multiple bond and pi bonds occur as the second or third in a multiple bond so for example if we'd like to know how many sigma bonds and how many pi bonds there are in this molecule let's begin counting as follows all single bonds are sigma i count six single bonds and the first of all multiple bonds so there i count one two for a total of eight sigma bonds pi bonds are found by taking what's left the second the second and third of multiple bonds so i count three pi bonds in that particular structure of the 11 total bonds 1 2 3 4 5 6 7 8 9 10 11. there were 11 bonds in that structure eight of those 11 are sigma the other three are considered pi all sigma bonds align on the same axes all pi bonds overlap either above or below the axes in this lesson we talk about hybridized atomic orbitals sp3 for example when we talk about hybridized orbitals what we're doing is promoting electrons from a ground state to an excited state let's kind of dive into this by considering methane ch4 we had drawn methane a while ago and it just had carbon as the central atom and four hydrogens attached to it if i write the ground configuration or carbon just to remind ourselves carbon is number six on the periodic table to write six electron configuration six electrons in its configuration 1s2 2s2 2p2 gives us 6 electrons knowing that these are the valence electrons those that are interesting to the chemist i have two electrons in the s of the second energy level and two more electrons in the p sublevel leaving one completely empty at this point when we examine this i only see two electrons available for bonding a bond is a shared pair so in order to share i have to have an unpartnered electron and right now this configuration would make it seem that carbon only has two electrons available for bonding and clearly we know that that's not the case because carbon forms four bonds so how does it achieve this when we write an excited state configuration we're going to promote one electron from the 2s orbital into the empty orbital of the carbon's p sub z so a moment ago i had both electrons in the 2s and i had 2 electrons in the 2 p sub z and pick that in x 2p sub x 2b some y 2 p sub z but i'm going to take this electron and promote it to the empty orbital of the two p sub z it's like giving everyone elbow room i have an empty bedroom and somebody's sharing in the 2s i'm going to allow this to just be promoted to that empty space and when i've done this i'm actually hybridizing hybridizing the orbitals to be degenerate remember what degenerate means it means of equal energy so instead of having a lower energy of s and higher energy of p i now have four equal energies of an s p three atomic orbital an sp3 is a hybridized orbital that was obtained by promoting the s electron into the empty p sub z now i want to make this clear a 2s energy level and a 2p energy level the hybridized lies between them i'm promoting the s but it's not quite as high as a two sub p i have four equal energy levels and now i can clearly see why carbon has room to form four bonds this then gets abbreviated as s sp3 all sp3 hybridized energy levels promote the electron from the s orbital up into the p orbital to create four of equal energy these are all tetrahedral in molecular geometry we have a central atom carbon four hydrogen attached back and the molecular geometry here for an sp3 hybridized orbital is called tetrahedral there are four electron domains an electron domain is either a bond just one word bond or lone pairs so here i can see one two three four my exponents have to add to four so s1 p3 gives us four electron domains we call that an sp3 hybridized orbital and bond angles are 109.5 let's look at another common hybridized orbital in organic chemistry called the sp2 hybridized orbital notice here s1p2 those exponents add to three so that means we have three domains an sp2 has three domains and a domain is either a bond or an electron pair in this carbon i see one two three domains now remember a domain is a bond or an electron pair the bond can be single double or triple it does not matter it still counts as one domain so when i look at ethylene this carbon has one two three domains now that tells me the sub the superscripts add to three an sp2 hybridization and this is the key sp2 in terms of its molecular geometry is known as trigonal planar it lies flat on a piece of paper all lies in the same plane this is so critical for organic mechanism that the carbon itself and the hydrogens leading out to the next carbon and its hydrogens are all in the same plane trigonal planar when we hybridize we promoted one of the s electrons into the p orbital and that gave us an sp2 hybridization and the last hybridization very common in organic is called an sp hybridized orbital notice the s and the p add to a total of two for their superscripts that tells me there's two domains notice this carbon in acetylene it has one two domains when i say the word domain it could be a single bond a double bond or a triple bond or a lone pair all of those count as a single domain so even though acetylene has a triple bond it still counts as one domain here is a single bond still counts as one domain so a total of two domains simply means that we have an s one p one one plus one gives us two for two domains this means that the molecular geometry for this molecule is linear with a bond angle of exactly 180 degrees just like a straight line this also is a 180 degree bond angle so it's a very linear oops i put the 0.8 180 degree bond angle so an sp3 is four domains it will always be tetrahedral that means the bond angle is about 109.5 think about its geometry and you have these like like the base of a pyramid with a straight up point and sp2 hybridized tells me that there's three domains its molecular geometry will be trigonal planar just like an equilateral triangle so just draw an equilateral triangle with the bond in the center and you've got that molecular geometry and these will all be 120 degrees so a flat triangle and of course sp tells me there's two domains and two domains means that it's a linear molecule with a bond angle of 180 degrees shall we try some what is the hybridization at each of these carbons and let's just look at this structure first let's number one two three four five six seven we have seven carbons in that particular structure carbon number one has one two domains and therefore we know that any compound here would be an sp hybridized orbital two domains carbon two also has two domains so it is also an sp hybridized orbital carbon three here's carbon three notice that it has four bonds so that means there's four domains that means that it's an sp3 hybridized carbon here's carbon four carbon four has one two three domains remember that even if it's a double bond it still counts as one domain so that's an sp2 trigonal planar 120 degree bond angle carbon five carbon five has one two three domains therefore we know it's an sp2 hybridized carbon remember even though it's a double bond it's still one domain carbon six carbon six has four single bonds so that's four domains that tells me that it's an sp3 hybridized carbon and the last has four single bonds telling me that four single bonds is four domains so that's an sp3 hybridized carbon let's try this carbon here i count one two three domains that's sp2 this carbon is also sp2 with three domains this carbon has three domains it is also sp2 and finally something different here that particular carbon has four domains so it's p3 hybridized it's kind of easier if you just practice here's a carbon with two domains so therefore it's linear with an sp hybridization and this carbon has three domains therefore it's an sp2 trigonal planar hybridized carbon let's talk briefly about bond strength and bond length in this particular slide i remember learning in general chemistry that a single bond is the longest bond and as i begin increasing electron density between the two carbons like in a double bond i decrease bond length all the way to a triple bond which is very electron rich between the two carbons and you have the shortest bond the longest bond is a single bond the shortest bond is the triple bond and the in in-between length of course would be the double notice that the amount of energy needed to break those bonds significantly increases with the more shared pairs between them to break this triple bond requires 820 kilojoules for every mole of energy to break this double bond between the carbons requires 632 kilojoules per mole and the least or easiest to break the least amount of energy 368 kj's per mole for a carbon to carbon single bond so the bond length decreases increasing number of bonds and to practice that simply rank the indicated bonds in terms of increasing length here we have a triple bond here we have a double bond and in this structure i have a single bond the single bond would be the longest and the shortest would be the triple and in between would be the double bond and that's easy enough to review i'm going to stop the video here and our last lesson will talk about the vesper theory and intermolecular forces two major topics left to go in our chapter