Summary: Study with Video Lessons, Practice Problems & Examples

Understanding the ground state electron configuration is crucial in organic chemistry, as it lays the foundation for concepts like electronegativity and the octet rule. Electrons fill orbitals according to the Aufbau principle, with configurations represented in both full and condensed forms. Electronegativity trends indicate that elements attract electrons more as you move right and up the periodic table. The octet rule emphasizes achieving eight electrons through bonding, while formal charges help verify Lewis structures. Recognizing functional groups, such as hydrocarbons and carbonyls, is essential for predicting reactivity in organic compounds.

Introduction to Organic Chemistry Channel

Hey, everyone. Welcome to the Organic Chemistry Channel in Pearson Plus. My name is Jules, and I'll be one of your guides through our content. Our job is to help you actually understand concepts from lectures so that you can ace your exams. I have my bachelor's degree in chemistry, and I've been helping college students in chemistry for nearly 15 years. And because of that, I'll be able to pass along insights and tips to help you better succeed and excel this semester.

In this organic chemistry course, we'll talk about common topics such as substitution versus elimination reactions, additional reactions of alkenes and alkynes, electrophilic aromatic substitution (EAS) reactions of benzene and its derivatives, nucleophilic acyl substitution (NAS) reactions of carboxylic acids and their derivatives, and so much more. We've developed our course with you in mind so you can use it along with your textbook.

So, Jules, how does this work? Well, you're going to watch our concept videos where we talk about these topics in detail. We're going to explain concepts, equations, and problem-solving techniques that you need to know. We always make sure to do real problems within our lessons, helping you to build up your skill and confidence so that you can eventually do these on your own.

Now, our lessons are built to be engaging. So you're not just going to be watching, you're going to have a downloadable copy that you can write along and follow along with me. That's right, I'm going to be there with you, working on examples, helping you, preparing you for your homework, and eventually for exam day. If you ever need additional help, feel free to reach out to the team by leaving a comment on the video, and we'll get back to you.

So again, welcome to Pearson Plus for organic chemistry. We're super excited to have you here. Whenever you're ready, let's get started.

Introduction to General Chemistry Review

Hey everyone. It's great to have you here on our organic chemistry channel on Pearson Plus. Now, before we begin, we're going to talk about a few topics from general chemistry 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 Electron Configuration

The ground state electron configuration is just the distribution of electrons, 1s, 2s, 2p, 3s, and so on, within orbitals using the Aufbau principle. Aufbau means "build up" in German. The Aufbau principle itself says, starting from 1s, electrons fill lower energy orbitals before moving to higher energy orbitals. 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 non-metals. So for the most part, you're not going to 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 1s1. Helium would be 1s2, because we're dealing with our second electron. We get to our third electron, the 1s2 is filled, so we move on to 2s 1. This will be 1s22s2, which is why we go to 2p1. Continuing, this will be 3s, so this will be 1s2 2s2 2p6 3s1, eventually getting us over here, where this would end with 3p1.

Looking at this pattern, let's see what we can do in terms of this example. We'd say write the ground state and condensed electron configurations for the following element, phosphorus. Here is phosphorus right here. It tells us its 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 1s2 2s2 2p6 3s23p1. And all we got to do is count to phosphorus. So that would be for the ground state, it'd be 1s2 2s2 2p6 3s2 3p3, because it's 3p1, 3p2, 3p3. 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 noble gas we pass before we get to phosphorus is neon. So neon is the element right here. So its condensed electron configuration would start out with the noble gas neon, and then we'd have 3s2 3p3. So this will represent our ground state and condensed electron configuration for the phosphorus atom.

Electronegativity

Electronegativity is going to be one of the most important periodic trends you need to remember when studying organic chemistry. Now, electronegativity itself is just the measurement of an element's ability to attract electrons to itself. And the periodic trend is that electronegativity 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 electronegativity increases. Now if we look at the periodic table here, we can see that a vast majority of it is greyed out. For example, here are our noble gases; we don't include them in electronegativity. They don't want to attract electrons to themselves because they're perfect. Inorganic 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 electronegative 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 electronegative group 7A element?" So, group 7A are our halogens. So we're looking at these here. If we look at our choices, sulfur and oxygen are not even in group 7A. So they would be out. So who's left? Bromine, iodine, and chlorine. Remember, as we head up a group, our electronegativity increases. Fluorine is not an option, so that's not going to be an answer. So the next one up would be chlorine. That means option D is the correct answer within this example question.

Octet Rule

Now the octet rule is the tendency of most main group elements in achieving 8 octet electrons by way of chemical bonding. Here we have valence electrons. These are just the electrons that an element possesses based on their group number. So if you're an element in group 3A, you'd have 3 valence electrons. Now, shared electrons are different. These are the electrons an element gains through a chemical bond. In organic chemistry, we're concerned with nonmetals, so we're concerned with covalent bonds. Because of this, remember, covalent bonds involve 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 following 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 about oxygen. Now remember, oxygen is a main group element. Oxygen wants to follow the octet rule, so ideally, you'd want 8 octet electrons. So that would mean that the answer is going to be either A or D. Now, its valence electron number is based on its group number. It is in group 6A, so it should have 6 valence electrons. Those are the electrons that are in red. These 6 electrons are what it comes to the table with when trying to form this molecule. And then it has 2 shared electrons. The shared electrons are the electrons it picks up by making covalent bonds with the hydrogen and the carbon. So, namely, this electron and this electron. So that gives it 2 electrons. Remember, when you add up your valence and your shared, it should give you back your octet electrons for most main group elements, giving us 8. That means, out of all the options, option B is the correct answer.

Formal Charge

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. Some key takeaways from this is that the only allowable formal charges for an element can be either negative 1, 0, or plus 1. 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 it might not be the best way to draw it. If you add up all the formal charges in your compound, that will equal the overall charge of the compound.

Now, when it comes to the formal charge formula, we're going to say that the formal charge equals valence electrons, which remember is just the group number of your element, minus your bonds that element is making, plus nonbonding electrons. Notice that this is in parentheses, so remember your orders of operation to calculate this correctly. And when it comes to nonbonding electrons, they are counted individually.

Here in this example, it says, "Determine the formal charges of each element within a thionyl cyanate ion." Nitrogen is in group 5A, so it has 5 valence electrons. Carbon would have 4. Sulfur would have 6. Coming back to nitrogen, we see it making 3 bonds with the carbon, and it has 2 nonbonding electrons. So, that would be 0 for its formal charge. Then we have carbon. It's in group 4A, so 4-4 bonds it's making total, and it has no nonbonding electrons, so it's also 0. Sulfur is in group 6A, so 6-1 bond it's making, and then if we count up each of these electrons individually, it adds up to 6. This would give a -1. So, we have 0 plus 0 plus minus 1 gives me a minus 1 overall charge for the thiocyanate ion, so this will be minus one. This makes sense because this is a polyatomic ion, which has a charge of minus 1. Right? So this is how we can utilize formal charge to see if a structure makes sense.

Drawing Lewis Dot Structures

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 2 chlorines in this molecule. Step 1 says determine the total number of valence electrons of the structure. Carbon is in group 4a, oxygen is in group 6a, and then we have 2 chlorines, each one in group 7a. So that's 4 + 6 + 14 which gives me 24 electrons. Now remember, your valence electrons are equal to your group number of electrons, so that's how we knew. Step 2 says, place the least electronegative 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 electronegative. And then halogens, they only make single bonds as a surrounding element. So here, carbon is the least electronegative, and it forms single bonds initially with its 3 surrounding elements. Now here we're going to add electrons to all the surrounding elements until they have 8 electrons, kind of following the octet rule. The exception here is that hydrogen only needs 2 electrons to be like helium, so it doesn't need 8. So if we do the math here, we're using 6 valence well, 6 electrons total so far. So there are 18 left. So that's 2, 4, 6, 8, 10, 12, 14, 16, 18. We have none left. So step 4, we can't place any remaining electrons in the central element, there's none left. Now, if any elements don't have 8 octet electrons, so 8 electrons around them, then we can add double or triple bonds. In step 6, 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 8 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 chlorines. Which means the double bond has to be formed between the carbon and the oxygen. So I'll erase one of these lone pairs here, and then I'm going to make a double bond from it. So those 2 electrons there, make this double bond. In that way, the oxygen has 8 valence electrons, and the carbon has 8 electrons as well. So, they both have 8 octet electrons, so they're both following the octet rule. So, this would be the correct structure for our molecule.

Resonance Structures

Now, resonance structures represent a set of two or more valid Lewis dot structures for polyatomic species possessing at least one pi bond. In a resonance structure, we have the movement of only electrons from either a pi bond or a lone pair. Here, we have the two resonance structures of the nitrite ion. If we look, we can see that the double bond or pi bond is on the left side here, but in the other structure, it's 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 pi bond has to get out of the way, and it will come over here to this oxygen. And that's why we have this lone pair here. This is just an intro to resonance structures; they are going to become more advanced than 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 lone pairs. In this particular example, we can say that the double-sided arrows are used to show that the resonance structures are equivalent to each other. That's because electrons are still being shared between 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 resonance structures. We're going to say to draw the resonance structures, we place a dotted line, anywhere a pi bond has been. So here and here. Now in this example, it says, determine the remaining resonance structures for the possible nitrate ion, not nitrite ion here, which is NO₃^-. So here, our pi 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 lone pairs. This oxygen up here would have three lone pairs. This one will have three lone pairs, and the overall charge is still -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 lone pairs. This one here has three, and this one here has three. And the overall charge is still the same. These use 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.

Hybridization

The hybridization of a central element can be connected to its number of electron groups. When I mention electron groups, I am referring to the lone pairs on the central element plus its bonding groups, which consist of its surrounding elements. In organic chemistry, we focus on nonmetals and adhere to the octet rule, simplifying the memory process for different types of hybridization. We typically consider up to 4 electron groups. When observing two electron groups, they would display an electron geometry of linear, and its hybridization would be sp. Imagine that there is a one beside s and p. Thus, it is represented as s1p1. If we add these up, 1+1, it indicates two electron groups on the central element. With sp hybridization, s orbitals present a single spherical shape, while p orbitals group into three dumbbell shapes. Therefore, sp suggests the involvement of these two hybridized orbitals, leaving two p orbitals unhybridized.

If you have three electron groups positioned trigonally planar, the hybridization will be . Omitting the invisible number one, it can be articulated as sp2, which translates to 1+2, amounting to three electron groups. This implies that the s orbital and two of the p orbitals are hybridized, leaving only one p orbital unhybridized. Finally, four electron groups describe a tetrahedral electron geometry. The hybridization in this case is sp3, calculated as 1+3, totaling four electron groups. In this instance, all orbitals are hybridized, thus none remain unhybridized.

Let us consider an example where we draw and determine the hybridization and number of unhybridized orbitals for a covalent compound such as HCN. The least electronegative element is typically placed in the center, which would be hydrogen, but because hydrogen cannot occupy this position, carbon is centralized. Hydrogen generally forms a single bond, thus requiring only two electrons to attain a similar electronic configuration as helium. Ideally, carbon forms four bonds, and nitrogen, belonging to group 5A with five valence electrons, prefers forming three bonds. Analysis of this structure reveals that carbon, as the central element, is surrounded by two electron groups - one hydrogen and one nitrogen. These two electron groups indicate that carbon's hybridization is sp, involving the hybridization of an s orbital and one p orbital, and leaving two p orbitals unhybridized. This analysis assists in conceptualizing hybridization related to various organic molecules.

Molecular Polarity

Now molecular polarity is the polarity that arises for an entire molecule. This is incredibly important because the polarity of a 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, a nonpolar molecule is any hydrocarbon. So if you have only carbons and hydrogens, you are nonpolar by default, and any 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 has 0 lone pairs. So if we take a look here, we are looking at linear, trigonal planar, and tetrahedral. These are perfect shapes. Here we only concern ourselves with electron groups 2, 3, 4, because in organic chemistry, we're looking at nonmetals. We're trying to adhere to the octet rule a lot of the times, so we need to go beyond these numbers of electron groups. The other shapes are not perfect because they possess at least one lone pair. Now here in this example it says, determine if nitrogen trifluoride is polar or nonpolar. Nitrogen is the least electronegative element here, so it goes in the center. Nitrogen is in group 5a, so it has 5 valence electrons. Fluorine is a halogen, and halogens make single bonds when they are surrounding elements. Fluorine itself has 7 valence electrons because it's in group 7a. So this is our structure. Now, it's not a hydrocarbon because it doesn't have only carbons and hydrogens so that is 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 0 lone pairs. So here we would say that it is a polar molecule. And it's as simple as that when you apply these rules about molecular polarity.

Functional Groups

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 the compound's reactivity. Here, we break up our functional groups into three categories: our hydrocarbons, those without carbonyls, and those with carbonyls. Let's look at hydrocarbons first. Now, hydrocarbons are just made up of hydrogen and carbon. If you have sp³ carbons connected to each other, and then single bonded to hydrogens, you're an alkane. An alkene is when you have at least two carbons double bonded to each other. An alkyne 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 carbonyl groups, remember: a carbonyl is when you have a carbon double bonded to oxygen. So here we're talking about the ones without carbonyl groups. Here, an alkyl halide is when we have an sp³ carbon connected to X. X here represents a halogen, elements from group 17. So fluorine, chlorine, bromine, or iodine. Next, we have an amine, that's when you have an sp³ 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 sp³ carbon connected to O. And ether is when you have COC. A thiol is when you have a carbon single bonded to an SH group.

Now with carbonyls, first we have aldehyde, which is a carbonyl connected to at least one hydrogen. A ketone is when you have a carbonyl connected to a carbon on either side. So, it must have a carbon on each side here. And acyl chloride is when we have a carbonyl connected to a Cl. An amide is when you have a carbonyl single bonded to a nitrogen. A carboxylic acid is a carbonyl group connected to an OH group, single bonded to it. Then finally, an ester is when we have a carbonyl, 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 and will start dealing with at some point in organic chemistry.

Here’s what students ask on this topic:

What is the Aufbau principle and how does it apply to electron configurations?

The Aufbau principle, from the German word 'aufbauen' meaning 'to build up,' states that electrons fill lower energy orbitals before moving to higher energy orbitals. This principle is crucial for determining the ground state electron configuration of an element. For example, in the case of phosphorus (atomic number 15), the electrons fill the 1s, 2s, 2p, 3s, and 3p orbitals in that order, resulting in the configuration 1s² 2s² 2p⁶ 3s² 3p³. This principle helps predict the chemical behavior of elements by understanding their electron arrangements.

How does electronegativity vary across the periodic table and why is it important in organic chemistry?

Electronegativity is a measure of an element's ability to attract electrons to itself. It increases as you move from left to right across a period and from bottom to top within a group on the periodic table. This trend is important in organic chemistry because it influences how atoms in a molecule share electrons, affecting molecular polarity and reactivity. For instance, in a molecule like H₂O, oxygen is more electronegative than hydrogen, leading to a polar molecule with significant implications for its chemical behavior.

What is the octet rule and how does it apply to chemical bonding?

The octet rule states that atoms tend to form bonds to achieve eight electrons in their valence shell, resembling the electron configuration of noble gases. This rule is particularly relevant for main group elements. For example, in a molecule like methane (CH₄), carbon shares electrons with four hydrogen atoms to complete its octet. Understanding the octet rule helps predict the bonding patterns and stability of molecules in organic chemistry.

How do you determine the formal charge of an atom in a molecule?

To determine the formal charge of an atom, use the formula: Formal Charge = Valence Electrons - (Bonds + Non-bonding Electrons). For example, in the thiocyanate ion (SCN^-), nitrogen (group 5A) has 5 valence electrons, makes 3 bonds, and has 2 non-bonding electrons. Thus, its formal charge is 5 - (3 + 2) = 0. Formal charges help verify the correctness of Lewis structures and ensure the overall charge of the molecule is accurate.

What are resonance structures and why are they important in organic chemistry?

Resonance structures are different Lewis structures for the same molecule that show the delocalization of electrons. They are important because they provide a more accurate depiction of the electron distribution within a molecule. For example, the nitrite ion (NO₂^-) has two resonance structures where the double bond can be on either oxygen. The true structure is a resonance hybrid, an average of these structures, which helps explain the molecule's stability and reactivity.