Summary

Hey, everyone. It's great to have you here on our Organic Chemistry channel on Pearson Plus. Now, before we begin, we're gonna talk about a few topics from Gen chem that carry over into organic chemistry. These topics will help to form the foundation for many of the theories you're going to see this semester. They'll give you a better understanding of the course overall and give you a better chance at success this semester. So whenever you're ready, let's begin.

The ground state electron configuration is just the distribution of electrons, one S two S two P three S and so on within orbitals. Using the Aha principle off means build up in German. And the AA principle itself says, starting from one S electrons fill lower energy orbitals before moving to higher energy orbitals. And when we talk about the condensed electron configuration, this is just a quicker way of writing it out here. We'd say we start at the last noble gas before the desired element. If we look at our periodic table, we break it up into different blocks. We have our S block, our P block, our D block and our F block organic chemistry is really concerned with our nonmetals. So for the most part, you're not gonna deal with D blocks and F blocks. And if we take a look here at this periodic table, we start out with hydrogen here, which is one S one helium would be one S two. Because we're dealing with our second electron, we get to our third electron, the one S is filled. So we'll move on to two S two S one. This would be one S 22 S two, which is why we go to two P one over here continuing, this will be three S so this will be one S 22, S 22 P six and three S one eventually getting us over here where this would end with three P one. Looking at this pattern. Let's see what we can do. In terms of this example, we'd say write the ground state and condense electron configurations for the following element phosphorus here is phosphorus right here I tell us if atomic number Z is 15, so it has 15 electrons. So following the pattern that we just saw in this spot here that I circled, this would be one S 22 s 22 P 63, S 23 P one and all we have to do is count to phosphorus. So that would be for the ground state, it be one S 22, S 22, P 63, s 23 P three because it's three P 13 P, 23 P three. If you add up all the electrons that we have here, it would add up to 15, which is equal to the atomic number. Now, the last double gas we passed before we get to phosphorus is neon. So neon is the element right here. So it's condensed electron configuration, we start out with the noble gas neon and then we'd have three s 23 P three. So this will represent our ground state and condense electron configuration for the phosphorus atom.

Electron negativity is gonna be one of the most important periodic trends. You need to remember when studying organic chemistry. Now, electron negativity itself is just a measurement of an element's ability to attract electrons to itself. And the periodic trend is that electron negativity increases, moving from left to right across a period and going up a group. So basically, as you're heading towards the top right corner of the periodic table, you would say that your electron negativity increases. Now, if we look at the periodic table here, we can see that a vast majority of it is grayed out. For example, here, our noble gasses, we don't include them in electron negativity. They don't want to attract electrons to themselves because they're perfect. And organic chemistry itself mainly focuses on nonmetals, which is why you can see a lot of the nonmetals here. Now, a few select metals are also included because you will see them pop up from time to time within different reactions. Most of the time they'll be spectator ions. So it's still important to know what their electron negative pattern will be. But just look at these numbers, learn the general trend memorizing numbers is not. What's most important. If we look at this example question, it says which of the following represents the most electron negative group seven A element. So group seven A are our halogens. So we're looking at these here, if we look at our choices, sulfur and oxygen are not even in group seven A, so they would be out. So who's left bromine, iodine and chlorine? Remember as we head up a group, our electron negativity increases, chlorine is not an option. So that's not gonna be an answer. So the next one up would be chlorine. I mean option D is the correct answer within this example question.

Now, the octet rule is the tendency of most main group elements in achieving eight octet electrons by way of chemical bonding. Now, here we have valence electrons. These are the just the electrons that element possesses based on their group number. So if you're an element in group three, a, you'd have three valence electrons. Now shared electrons are different. These are the electrons that element gains through a chemical bond in organic chemistry. We're concerned with nonmetal. So we're concerned with covalent bonds because of this. Remember, covalent bonds are the sharing of electrons. So this is where shared electrons come into play with this information. We can say that octet electrons equal your valence electrons plus your shared electrons. If we take a look here at this example, it says which of the following statements is true in terms of the falling compound. Now, this compound itself is called methanol. Don't worry about the name. That's just what it's called. But let's focus on it here. It says a lot of things about oxygen. Now remember oxygen is a main group element. Oxygen wants to follow the octet rule. So ideally you'd want eight octet electrons. So that would mean that the answers can be either A or D. Now its valence electron number is based on its group number. It is in group six A. So it should have six valence electrons. Those are the electrons that are in red, these six electrons are what it comes to the table with when trying to form this molecule. And then it has two shared electrons. The shared electrons are the electronic picks up by making covalent bonds with the hydrogen and the carbon. So namely this electron and this electron, so that gives it two electrons. Remember when you add up your valence and you're shared, it should give you back your oc 10 electrons for most main group elements. So giving us eight, that means out of all the options. Option B is the correct answer.

Now, sometimes you might draw an organic molecule and you aren't quite sure if it's drawn correctly. Well, that's where formal charge can help. It's used to determine and check to see if you drew your Lewis dot Structure correctly. Now some key takeaways from this is that the only allowable formal charges for an element can be either negative 10 or plus one. So if you took the formal charges of all the elements within your molecule, and they don't give you one of these three numbers that just means that that might not be the best way to draw it. Now, if you add up all the formal charges in your compound, well, that will equal the overall charge of the compound. Now, when it comes to the formal charge formula, we're gonna say that formal charge equals valence electrons, which remember is just the group number of your element minus your bonds that element is making plus non-binding electrons. Now notice that this is in parentheses. So remember your orders of operation and calculate this correctly. And when it comes to nonbinding electrons, they they are counted individually. Now, here in this example, it says, determine the formal charges of each element with a thin cyanide ion. Now, nitrogen is in group five A. So it has five valent carbon would have four sulfur would have six coming back to nitrogen. We see it making 123 bonds with the carbon and it has 12 non-binding electrons. So that would be zero for its formal charge. Then we have carbon, it's a group four A or four minus the four bonds, it's making total and it has no non bonding electrons. So it's also zero sulfur is in group six A. So six minus the one bond it's making. And then if we count up each of these electrons individually, it adds up to six. This will give me minus one. So we have zero plus zero plus minus one, gives me a minus one overall charge for the thiocyanate ion. So this will be minus one. This makes sense because this is a poly atomic ion which has a charge of minus one, right. So this is how we can utilize full charge to see if a structure makes sense.

Now many possible Lewis dot structures exist but there are rules to draw the best structure here. In this example, it says draw the Lewis dot structure for the following molecule. So we have carbon oxygen and two chlorine in this molecule. Step one has determined the total number of valence electrons of the structure. Carbon is a group four A oxygen is in group six A. And then we have two coins each one in group seven A. So that's four plus six plus 14 gives me 24 electrons. Now remember your valence electrons is equal to your group number of electrons. So that's how we know step two says place the least electron negative element in the center and connect all elements with single bonds. Now, here there are exceptions. First, hydrogen never goes in the center even if it's the least electron negative and then halogens they only make single bonds as a surrounding element. So here carbon is the least electron negative and it forms single bonds initially with its three surrounding elements. Now, here we're gonna add electrons to all the surrounding elements until they have eight electrons. So kind of following the octet rule exception here, hydrogen only needs two electrons to be like helium. So it doesn't need eight. So if we do the map here, we're using uh six valence elect well, six electrons total so far. So there's 18 left. So that's 2468, 10, 12, 14, 16, 18, we have none left. So step four, we can't place any remaining electrons in the central element. There's none left. Now, if any elements don't have eight Octa electrons or eight electrons around them, then we can add double or triple bonds. In step six. We only use that if we're still not sure if the structure is drawn correctly, we can utilize formal charge to check and see if it's drawn correctly. Now, if we take a look here, the surrounding elements all have eight electrons around them but not the carbon in the center. Now, this part here says that halogens can only make a single bond. So I can't make a double bond with either one of those chlorine, which means the double bond has to be formed between the carbon and the oxygen. So I era one of these alone pairs here and then I'm gonna make a double bond from it. So those two electrons there make this double bond. And in that way, the oxygen has eight valence electrons and the carbon has eight electrons as well. So they, they both have eight Octa electrons. So they're both following the octet rule. So this would be the correct structure for our molecule.

Now, resonant structures represent a set of two or more valid le dot structures for polytonic species possessing at least one pie bond. Now, in a resonance structure, we have the movement of only electrons from either a Python or lone pair. Here, we have the two resonance structures of the nitrite ion. If we look, we can see that the double bond or pyon is on the left side here, but in the other structures on the right. So how exactly does this occur? Well, this is what happens we have the movement of a lone pair from the oxygen on the right side coming here to make a double bond that helps to make this double bond here. But moving that electron over means that nitrogen will be breaking the octet rule. So what has to happen is that this Pie Bon has to get out of the way and it will come over here to this oxygen. And that's why we have this loan pair here. Now, this is just an intro to resonance structures. They're gonna become more advanced in this but just realize that this is the fundamental approach when it comes to resonance structures. It's the movement of only electrons from pi bonds and or alone pairs. In this particular example, we can say that the double sided arrows are used to show that the resonance structures are equivalent with each other. And that's because electrons are still being shared between a nitrogen and oxygen in both structures. That's not always going to be the case in this particular one, it is. So they're of equal or similar energy. Now, the real structure is represented by the composite or average of the resonance structures and that's called the resonance hybrid. Now, the resonance hybrid is a composite of all the major uh resonance structures. We're gonna say to draw the resonance structures, we place a dotted line anywhere a pie bond has been. So here and here. Now, in this example, it says, determine the remaining resonance structures for the possible nitrate ion nitrate ion here uh which is no three minus. So here our pie bond is with this oxygen up here, but it could easily be with any of the other oxygens. So here we could have the double bond being here instead. And now it has two long pairs. This oxygen up here would have three lone pairs. This one will have three long pairs and the overall charge is still minus one or we could have this as our resonance structure where the double bond is on the oxygen on the bottom left, it has two loan pairs. This one here has three and this one here has three and the overall charge is still the same. These are the same type of double sided arrow because we still have the movement of electrons between nitrogen and oxygen, they're all equivalent to one another, right. So this is just the beginnings of our approach to resonance structure. So keep that in mind as you delve deeper and deeper into organic chemistry.

The hybridization of a central element can be connected to its number of electron groups. Now, when I say electron groups, I'm talking about the lone pairs on the central element plus its bonding groups. Now, bonding groups is just its surrounding elements. Now, in organic chemistry, we're concerned with nonmetals and we're also concerned with the octet rule. So that makes it much easier to remember the different types of hybridization. We only don't have to go up to four electron groups. Now, if we take a look here, two electron groups would have an electron geometry of linear, its hybridization would be sp now just imagine that there's a one here and the one here there, we don't see them but assume they're there. So it's s one P one, if we add them up one plus one, that gives us two, that tells us how many electron groups that sensor element has on it. Now, if you're sp hybridized, remember s orbitals have one shape of sphere but P orbitals, there are three of them, three dumbbells sp would just mean that we're talking about these two, these would be our two hybridized orbitals, which would mean we'd have left these two p un hybridized orbitals. Now, if your three electron groups are trigonal planar, so here your hybridization will be SP two. There's a one without that we don't see. So here we'd say one plus two, gives me 33 electron groups. Sp two means that the S orbital and two of the P orbitals are hybridized, meaning that there's just one P left that's un hybridized. Finally, if you have four electron groups, you are tetrahedral. In terms of your electron geometry, your hybridization is sp three, that'd be one plus three, which is 44 electron groups around your central element. Here, everything is hybridized. So here you have nothing left un hybridized. Now, here, if we take a look at the example, it says, draw and determined the hybridization and number of un hybridized orbitals for the following covalent compound. So here we have HCN. So the least electron negative element goes in the center, which would technically be hydrogen. But remember, hydrogen can't go in the middle. So carbon has to go in the middle hydrogen only makes single bonds since it only needs two electrons to have the same number of electrons as helium carbon ideally wants to make four bonds. And that's OK because nitrogen would like to make three bonds. Nitrogen group, five aid has five valence electrons. This will represent our structure. If we take a look here, we'd say that carbon is our central element. It has two electron groups, this hydrogen and this nitrogen, two electron groups would mean that its hybridization is sp so it's s orbital is hybridized and one of its p orbitals is hybridized, meaning that we have two P orbitals that are left on hybridized. OK. So this is the way we think about hybridization when it comes to our different types of organic molecules.

Now, molecular polarity is the polarity that arises for an entire molecule. This is incredibly important because the polarity molecule can dictate the types of reactions that are possible with it when we study organic chemistry. Now here, what classifies as a nonpolar molecule? Well, we're gonna say nonpolar molecule is any hydrocarbon. So if you have only carbon and hydrogens, you're nonpolar by default and any non non hydrocarbon with a perfect shape. Now a perfect shape is where you have your central element and it has the same surrounding elements and it has zero on pairs. So if we take a look here, we're looking at linear trigonal planar and tetrahedral. These are perfect shapes here. We only concern ourselves with electron groups 23 and four because in organic chemistry, we're looking at nonmetals, we're trying to adhere to tit rules a lot of the time. So we don't need to go beyond these number of electron groups. The other shades are not perfect because they possess at least one loan pair. Now, here in this example, it says, determine if nitrogen tri fluoride is polar or nonpolar. Nitrogen is the least electron negative element here. So it goes in the center nitrogen is in group five A. So it has five valence electrons. Fluorine is a halogen and halogens make single bonds when they're surrounding elements. Fluorine itself has seven valence electrons consists in group seven A. So this is our structure. Now, it's not a hydrocarbon because it doesn't have only carbons and hydrogen. So that's out and it's not a perfect shape because although it has the same surrounding elements, it possesses a lone pair, you need to have zero lone pairs. So here we'd say that it is a polar molecule and it's as simple as that when you apply these rules about molecular polarity.

The vast majority of organic chemistry is basically the reactions of different functional groups. Now, a functional group is the part of a molecule that is recognizable and responsible for compounds reactivity. Here, we break up our functional groups into three categories are hydrocarbons, those without Carbonis and those with Carbonis. Let's look at hydrocarbons first. Now, hydrocarbons are just made up of hydrogen and carbon. If you are just sp three carbons connected to each other and then single bonded to hydrogens, you're an al cane. And now keen is when you have at least two carbons double bonded to each other. And alkin is when you have at least two carbons triple bonded to each other. This last structure here is a special type of ring called a benzene ring. You'll learn more about that. The further into organic chemistry you go now when it comes to Carboni groups, remember a Carboni is when you have a carbon double boded to an oxygen. So here we're talking about the ones without carbon yel groups. Here. An Alki hali is when we have an P three carbon connected to XX. Here represents a halogen elements from group at 47 A so fluorine, chlorine, bromine or iodine. Next, we have an amine. That's when you have an sp three carbon connected to a nitrogen, that nitrogen could have hydrogens connected to it or could have additional carbons as well. Alcohol. This is when you have an SP three carbon connected to oh and ether is when you have coc A thol is when you have a carbon single body to an sh group. Now, with Carbonis first, we have aldehyde which is a carbonyl connected to at least one hydrogen. A key tone is when you have a carbon e connected to a carbon on either side. So it has to have a carbon on each side here. And acid chloride is when we have a carbonyl connected to AC L an amide or am was when you have a carbon not connected single boded to a nitrogen. A carbolic acid is a carbonyl group connected to an oh group, single boded to it. Then finally, in a, when we have a Carboni single bonded to an oxygen which is then single bonded to another carbon. These represent the majority of the functional groups that you need to remember you're gonna start dealing with at some point in organic chemistry.