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Week 1 and 2 Notes

by: iceskatercjudd Notetaker

Week 1 and 2 Notes Chem 2310

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Organic Chemistry 1
Holly Sebahor
Class Notes
Organic Chemistry




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This 21 page Class Notes was uploaded by iceskatercjudd Notetaker on Tuesday September 6, 2016. The Class Notes belongs to Chem 2310 at University of Utah taught by Holly Sebahor in Fall 2016. Since its upload, it has received 186 views. For similar materials see Organic Chemistry 1 in Chemistry at University of Utah.

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Date Created: 09/06/16
Colby Judd Notes Summary Introduction: Welcome to my notes section for the Organic Chemistry Course. For each of my courses, I do notes in a different way that I deem to fit the needs of the course. For this class I will be taking three different styles of notes. My notes are uploaded weekly. The first style will be the reading notes. These notes will not include examples and will attempt to talk through the reading and make sense of it. It will pull out what I deem to be most relevant to the class and will offer hints that I find useful. Understand that these notes will be the most difficult to understand, as they do not include examples and are my first exposure to the subject material. I will create reading notes for both Klein and the Organic Chemistry Textbook. I will also make flash cards of certain molecule types and anything I find useful. The second style will be lecture notes. During lecture, I take what I call "active notes". This means that I spend most of my time in lecture paying attention to the professor and will only write down material I find difficult or relevant, I will not copy down the entire lecture. For ideas that I struggle more with, I write them down and will research them after class more or during a TA section. Because these notes are not easy to understand, I try my best to lay them out in an easy to read manner that is succinct after class. The third style will be my weekly summary. The weekly summary will come in a similar format as lecture notes and I will attempt to narrow down the most relevant information to two pages each week. Understand that I am taking this class for my first time as well, so there will be mistakes in my notes and ideas may not be expressed clearly. I attempt to correct those mistakes in my summary page as I understand the material best at that point. Using my notes may or may not help you depending on your learning style. I suggest as well that you utilize the summary page at least for my notes but think all the info. Is relevant to the course. Do the suggested coursework yourself; this is just for reference! If you have suggestions on the format of my notes to improve them, you may email me at Table of Contents: 1. Chapter 1 Reading 2. Klein Reading 1, 2, and 4 3. Lecture #2 Notes 4. Chapter 2 Reading 5. Klein Reading 3 6. Lecture #4 Notes 7. Chapter 3 Reading 8. Klein Chapter 8 (Postponed to the following week) 9. Summary ( flashcard) 08/23/2016 Organic Chemistry Chapter 1 Pre-lecture Structure and Bonding Understand the basic fundamentals of the periodic table in chemistry and how that influences chemical behavior of molecules. This should just be a review of general chemistry.  When atoms bond they become more stable and have lower energies  Constitutional isomers have the same molecular formula but are connected differently, creating new compounds  H, B, Be, P, and S are all common exceptions to the octet rule  When we draw resonance structures in chemistry, we are talking about compounds with the same placements of atoms but a different arrangement of electrons. Both resonance structures are not conditions of the compounds but rather a hybrid of them displays where the electrons exist, meaning that the electrons are delocalized between spaces. A molecule with two or more resonance structures is resonance stabilized.  When dealing with resonance structures, double bonds can turn into single bonds with an extra electron pair. Use curved arrows to display this change in resonance structures.  Heteroatom: Not carbon or hydrogen  A resonance hybrid structure is a composite of all the resonance structures and delocalizes electron density over a larger volume. Major contributors are better resonance structures to the hybrid, rather than minor contributors. This means that they have move bonds and fewer charges.  Hybrid resonance structures will represent one double bond with a dashed line, to show that the double bond moves between electron pairs and bonds. It will use partial charges as well to designate where charges move places.  Bond length increases as you move left and down on the periodic table.  Review basic VSEPR principles; o Groups refers to electron pairs and bonds o 2 groups = linear = 180 degrees o 3 groups = trigonal planar = 120 degrees o 4 groups = tetrahedral = 109.5 degrees  When drawing Lewis Structures, draw hydrogen bonds in front with a black wedge and those in back with dashed wedges.  Two types of organic structures, condensed structures and skeletal structures  Condensed structures represent the molecules involves by pairing similar molecules together and using their letter representation. Bonds are not drawn  Skeletal structures are for rings and chains of atoms. They show the skeletal structures and do not display carbon atoms. Carbon atoms are assumed to be at the end of bonds. Hydrogens are assumed to make carbon atoms tetravalent as well and are not displayed either. Heteroatoms display bonded hydrogen atoms.  When carbon atoms are charged it creates lone pairs. Don’ t draw lone pairs on heteroatoms.  Hybridization is the combination of multiple atomic orbitals to create the same number hybrid orbitals with the same shape and energy. Creates stronger bonds.  2 groups create 2 sp orbitals  3 groups create 3 sp^2 orbitals  4 groups create 4 sp^3 orbitals  IMPORTANT: Ethane is CH3CH3, all carbons are tetrahedral and sp^3 hybridized.  IMPORTANT: Ethylene is C2H4, all carbons are trigonal planar and each carbon is sp^2 hybridized. Central double bond makes rotating energy difficult.  IMPORTANT: Acetylene – C2H2, all carbon is linear, and each carbon is sp hybridized. Triple bond.  Pi bonds form on a different plane then sigma bonds making them weaker.  The number of electrons between two atoms, increases the bond strength, and shortens the bond length.  More s character, which is the amount of s relative to p, which means greater bond strength.  Bonds are polar when electrons are unequally shared. Main Takeaway: There are a variety of rules for the structures of molecules and compounds in chemistry. Remember that the position of electrons and bonds will influence the shape and strength of structures in chemistry. There are a few different ways to display these structures that either display a structural representation or molecular representation. Finally, resonance structures do not mean that electrons can occupy these two different positions, but rather give an idea of two extreme positions. Electrons are throughout the orbitals. Klein Organic Chemistry Chapters 1, 2, 4 08/29/2016 Chapter 1: This is one of the most basic chapters of the book and goes over some info. That you will need throughout this class. The first thing to know is that you represent carbons in line drawings as intersections of lines, rather than writing out carbon. You also assume that carbons use hydrogen atoms to create a full octet shell around it, unless it has a formal charge which will indicate a lone pair of elections. Chapter 2: This is also fairly basic, but there are a few rules to make sure of. First of all, a resonance structure is a movement of electrons in a molecule, not the movement of atoms which is an isomer. Resonance structures represent different electron placements within one molecule, but it is important to note that the actual structure is a hybrid of resonance structures. We use curved lines to represent the movement of electrons from one resonance structure to another, the tail is where they come from and the head where they go to. The tail must always come from a spot with electrons as well. There are two commandments that must be followed for drawing resonance structures. 1. Never break a single bond 2. Do not exceed an octet for second-row elements. When we do create resonance structures, we must always remember to write down the formal charge as well for the changes. Molecules cannot gain or lose charge as resonance structures, so if you begin with a net negative charge, you must end with a net negative charge on some other atom. This is important so we know the electronegativity of the molecule and the likelihood for electrons to be distributed in a certain area. There are other things that we should consider as well for resonance structures, when preparing to draw them. 1. Can we convert lone pairs into pi bonds? 2. Can we convert pi bonds into lone pairs? 3. Can we convert pi bonds into pi bonds? These are good questions to ask to start determining ways in which resonance structures can be formed. The book also mentioned 5 patterns that are helpful to remember when looking for resonance structures. 1. A lone pair next to a pi bond 2. A lone pair next to C^+ 3. A pi bond next to C^+ 4. A pi bond between two atoms, where one of those atoms is electronegative 5. Pi bonds going all the way around the ring. These five clues all suggest potential resonance structures based around these atoms. The final part of resonance structures is determining the major and minor contributing structures. Some resonance structures are more prominent in the total hybrid than others, this has to do with stability. There are four rules that can be followed to help determine what resonance structures contribute more to the final hybrid then others. 1. The most important is the resonant structure with the most covalent bonds. 2. The structure with the fewer formal charges is also important. 3. A structure with a negative charge on the more electronegative element will contribute more, the opposite is true if it is a positive charge. 4. Resonance structures that are equal are equivalent. Chapter 4: We will only deal with two types of orbitals, s orbitals and p orbitals. There can only be one s orbital around an atom, so after that has been exhausted we move on to p orbitals. You determine the number of orbitals by summing the number of atoms surrounding an atom along with the electron pairs. When you have 4 orbitals – you get 4 sp^3 orbitals. 3 is 3 sp^2 orbitals and so on. There are a few exceptions to this rule, but for now we will ignore those exceptions in the class. This is the hybridization of the atom and is used to determine the geometry of the atom as well. The geometries of the atoms follow the VSEPR model and are very simple: Sp^3 Tetrahedral 109.5 degrees 4 Bonds Sp^3 Trigonal Pyramidal <109.5 degrees 3 Bonds 1 electron pair Sp^3 Bent <109.5 degrees 2 Bonds 2 electron pair Sp^2 Trigonal Planar 120 degrees 3 Bonds Sp^2 Bent <120 degrees 2 Bonds 1 electron pair Sp Linear 180 degrees 2 bonds One thing that is important to know, is a lone pair that participates in resonance must occupy a p orbital. So if you have a Nitrogen with three bonds and an electron pair it would be sp^3 hybridized, unless that electron pair participates in resonance and forms a double bond which will result in sp^2 hybridized. 09/03/16 Organic Chemistry Chapter 2 Reading Chapter 2: Acids and Bases 2.1 The chapter begins with the Bronsted-Lowry definition of acids and bases, which will be the most common definition of Acids and Bases for this cours.e A Bronsted-Lowry (BL) Acid is a proton donor, in other words it has a hydrogen atom that it gives away. A BL Base is a proton acceptor; or as the professor prefers a "Hydrogen Attacker". It can make this bond with hydrogen by using lone pairs of electrons or breaking pi bonds. Breaking pi bonds requires a lot of energy, so is less likely than lone pairs picking up hydrogen. HA – Bronsted – Lowry Acid; B = Bronsted – Lowry Base The ions that are part of the base to balance the charge, cations known as salts. They are inconsequential and are known as counterions or spectator ions. 2.2 We call a reaction between a BL Acid and base, a proton transfer reaction and it involves the movement of hydrogen from the acid to the base. When documenting the reaction, we show the movement of electrons that result in the transfer of Hydrogen, not the movement of Hydrogen itself. So an electron pair typically will move from the Base to the hydrogen and the bond between hydrogen and the other molecules of the acid will break, where the electrons will move to the other molecules. This is more detailed in the class lecture notes and the summary notes at the back of this note packet. When an acid loses its proton, hydrogen, in a reaction, it is known as a conjugate base. A conjugate acid will be the result of gaining a proton. Another important note is the net charge must be the same before and after the reaction. Another way to think about acid-base reactions is as follows: Electron-rich species (bases) react with electron-deficient species (Hydrogens) 2.3 The strength of an acid is determined by how readily it will donate a proton. We use the value of pK_a to help determine the strength of an acid. pK_a is really a logarithmically simplified version of Ka so we are dealing with smaller and easier to utilize numbers. The important thing to note is the stronger the acid, the more the equilibrium shifts to the right of the reaction equation, and the smaller the value of pK_a. pK = - log K a a Two Rules of Acid Strength: 1. Strong Acid = Weak Conjugate Base 2. Strong Base = Weak Conjugate Acid 2.4 When looking at the equilibrium of acid-base reactions. Equilibrium will always favor the formation the formation of a weaker acid and base, so determine the pKa of the acids on both sides of an equilibrium equation and the equilibrium will be towards the higher value side. The base is only strong enough to deprotonate an acid, remove a proton from an acid, when the conjugate acid is weaker than the starting acid. 2.5 There are four factors that help adjust acid strength and they proceed in the order that I will discuss them in. When determining acid strength, you must follow this order as the first factors produce a larger change in acid strength than the last factors. It is also very important that you follow the SAME procedure when determining these factors acids. 1. Draw the conjugate bases of the acids 2. Use the factors to determine which conjugate base is the most stable 3. The most stable conjugate, has the most acidic acid Factor 1: The periodic table trend of the elements. Moving across the periodic table, acidity will increase as the electronegativity increases. This will be in the rightward direction. This is because, when a conjugate base is formed, the new electron pair will be more stable with an atom of higher electronegativity to hold the electron pair close. The more stable the electron pair with its central atom, the easier for the sigma bond between hydrogen and the central atom to break. The second part of this factor is the downward trend of the periodic table. As you move down along the periodic table, atom size will increase. The increase in atom size is a much stronger indicator of stability, overpowering the factor of electronegativity decreasing down the periodic table. This is because with more size, you can disperse charge over a larger volume on the atom and thus makes it more stable. MOST IMPORTANT: Acidity and conjugate base stability increase as you move to the right and bottom of a periodic table. Factor 2: Inductive effects is all about how charge can be dispersed across a compound. When you place a highly electronegative atom on the tail end of a compound, such as Fluorine, it will attract negative charge. This is perfect because a central atom, such as oxygen, utilizes its net negative charge to keep hydrogen, which is electron deficient attached. Fluorine is able to pull some of this negative charge from the oxygen through the sigma bonds in the compound, to result in a weaker bond between hydrogen and oxygen and a much more stable conjugate base. This will increase acidity. We call this process the electron-withdrawing inductive effect. Factor 3: Going back to the theory of Resonance, resonance structures resemble how delocalized electrons can be in a compound, over different areas of electron density. That means the greater number of resonance structures, the weaker the electron density throughout the compound. In terms of acidity and stable conjugate bases, this will increase stability and acidity. The more dispersed the electrons over a compound, the less likely the compound will be able to keep a strong hold on a proton. So as the amount of resonance stability increases in a conjugate base, the greater the acidity. Factor 4: Lastly, hybridization plays a key role in determining acid strength. As you increase the s-character of a compound, you are increasing the amount of low energy s bonds within a compound in comparison to high energy p bonds. In the long run, this means it will be easier to separate protons from an acid and will create more stable conjugate bases. It is important to note that the hybridization produces the smallest effect of the four factors on acidity and conjugate base strength. There is one last note to make for this section which is you can use these procedures to determine the hydrogen atom in an acid that is most likely to separate from the structure, by using these factors to determine the type of bonds it has. 2.6 This is very basic and it would be useful to make flashcards of the listed acids and bases below. I will also add a link to flash cards to the tail end of the notes section. Common Acids: 1. HCL 2. H 2O 4 3. Acetic Acid 4. TsOH Common Bases: - 1. OH 2. NH 2 3. H - 4. Carbanions – Negative charged carbon atoms 5. Triethylamine 6. Pyridine 2.7 This chapter contained information that was more interesting and applicable to real life. In the case of Aspirin, it is a member of a compound group called salicylates. Aspirin is useful for many medical reasons, but one of the great aspects of aspirin is the fact that it is able to deprotonate within different environments of our body. This means it will switch from its acidic form, called its neutral form, to its ionic form, which is its conjugate base. Within our body, the basic environment of our small intestine will deprotonate it to its ionic form, where it can travel through the bloodstream. From there it will be re-protonated to enter our cells and perform work. While this is mostly interesting information, it is likely useful to understand that aspirin is an acid and that property makes it useful for medical applications. 2.8 The final section of the notes is dedicated to Lewis Acids and Bases! It is important to know that a Lewis Acid is an electron pair acceptor. It is NOT A PROTON DONOR, meaning that some Lewis acids will be without hydrogen molecules, because they don't take electron pairs to give off hydrogens. They take electron pairs to form bonds with Lewis Bases. On the other hand a Lewis Base is an electron pair donor. BL and Lewis Bases are the same, but while all BL acids are Lewis Acids, the reverse is not true. Electron-rich species react with electron-poor species. Chapter 3 Reading Notes 09/05/16 Intro to Organic Molecules and Functional Groups 3.1 Functional groups are what distinguish different organic molecules. They are the heteroatoms and pi bonds within molecules and are the reactive part of the molecule. They distinguish the chemical properties and create electron-deficient sites on carbon. The carbon and hydrogen bonds of an organic molecule are mostly irrelevant, as far as chemical properties are considered. They are the skeleton of the molecule and abbreviated with the letter R often. 3.2 There are different types of functional groups. The first type are hydrocarbons. These functional groups only contain hydrogen and carbon bonds and are either aliphatic or aromatic. Aliphatic hydrocarbons can be alkanes (only single bonds and no functional sites), alkenes (Have a double carbon bond as a functional group), or alkynes (Have a triple carbon bond as a functional group). *An easy way to remember these namesndis the order of the second vowel. A comes first and has 1 bond, e comes 2 and has two bonds, y comes third and has three bonds.* Aromatic hydrocarbons have strong orders and are variations of benzene rings. The functional group for them will be called the phenyl group. We label the carbons in compounds depending on the number of carbons o atoms that are bonded to a certain carbon atom. A primary carbon (1 carbon) is bonded to one other C atom, a secondary to 2, tertiary to 3, and quaternary to 4. We also can label hydrogen atoms similarly, be giving them classifications depending on how many carbon atoms are connected to the carbon atom they are bonded two. primary for one carbon, secondary for two carbons, and tertiary for three carbons. Use the X notation. Another type of functional groups are groups that contain a carbon bonded to a electronegative heteroatom using a sigma bond. The polar bond will make carbon electron deficient. There are a long list of different functional groups that fit this description and will be included in the flashcards and various atoms within these molecules can be identified using the same process described above for hydrocarbons. The last type of functional group is a carbon bonded to an electronegative heterotam using a pi bond. Most commonly this will be a carbonyl group (CO). Labeling proceeds similarly except for amides, where you will count the number of carbon atoms bonded to the nitrogen bonded to the carbonyl group. Flash cards will better describe this. 3.3 Intermolecular forces are determined by functional groups. The strongest intermolecular forces are ionic bonds between oppositely charged particles. Will not expand as this is covered in general chemistry. Covalent compounds depend on intermolecular forces and functional groups. There are three types of forces. 1. Van der Waals Forces (London Forces): Depend on the electron densitry of a molecule. Because electrons move position, temporary dipoles are created and results in a weak interaction between two molecules. The larger the surface area will increase this effect as electrons are less closely held and more free to move (polarizability). The weakest of the three. 2. Dipole-Dipole Interactions: Forces between the dipoles of molecules that are created by the balance of positive and negative charges. 3. Hydrogen Bonding: The interaction between a Hydrogen atom bonded to O, N, or F to a lone pair on a O, N, or F atom in another molecule. This is the strongest and is common in water. 3.4 These forces alter the physical properties of molecules. The stronger the intermolecular forces, the higher the boiling point as it becomes more difficult to separate molecules. You can use a distillation lab technique to separate different liquids, by boiling the liquids. The more volatile liquid will boil sooner and evacuate to a receiver flask, leaving the other liquid behind. The melting point is higher for stronger IMF forces as well for the same reasons. Symmetry of a molecule does play another role in the melting point however, as a symmetrical molecule will pack better and be more difficult to separate in the melting process, then an asymmetrical molecule. Solubility is dependent on likes dissolving likes. So organic compounds dissolve organic compounds. Also, an organic molecule is only dissolvable in water if it has one polar functional group that can hydrogen bond, per 5 carbon atoms it contains. (This section discussed that ionic compounds are dissolved in an excess of water, to create an extreme amount of ion-dipole reactions. Likely not important currently.) Applications: *This section serves as less relevance to the material in class, but has applications elsewhere. Understanding the basic properties of chemistry discussed within these notes, should make understanding the problems possible* The next three chapters talk about applications of solubility in organic compounds. Some vitamins are water soluble so they can be used quickly within the body while others are fat soluble, so they can be used over extended periods of time and stored. Vitamin A is fat soluble and Vitamin C is water soluble. For soap you have a functional group attached to a long hydrocarbon chain. This makes one end hydrophilic and the other end hydrophobic. Allows soap to form micelles which are spherical shells that surround nonpolar hydrocarbons (AKA dirt) and remove them. Again we see the application in the cell membrane, which is composed of phospholipids with two hydrocarbon tails and an ionic head. The same idea forms as the ionic head will interact with water and the nonpolar tails will form an insoluble barrier from the outside of the cell. Uncharged organic molecules are able to easily cross this membrane. Ionophores are created to transport molecules across the membrane. These are complex cations that behave like a hydrocarbon externally but have oxygen atoms on the interior to bind special cations. The nonpolar hydrocarbon shell makes it possible to move through the barrier. 3.8 The final section will play off of Chapter 8 of Klein a bit more, but we will go over some basics for now. The first thing to note is that an electronegative heteroatom will pull charge and make a carbon atom electrophilic. A lone pair on a heteroatom will make it basic and nucleophilic. Pi bonds also create nucleophilic sites and are easily broken down. What does this mean exactly? - A nucleophile (Nu ) is an atom that likes a nucleus, positive charge. So an electron-deficient carbon (positive carbon) will react with a nucleophile and an electrophile (E ) likes a negative charge. An electron-rich carbon (negative carbon) will react with an electrophile. Klein Chapter 3 09/05/16 Acid Base Reactions Qualitative Assessment of Acidity The main thing to understand for this chapter is electron density and how that plays a role in the willingness of an acid to give up its proton (Hydrogen). This is determined by the stability of the conjugate base. The more stable the conjugate base, the more likely it will give up its proton as it isn't necessary to stabilize the electrons. The first factor in determining this is the atom that the charge is one. The more electronegative an atom is, the better it holds charge and the more stable the electrons. The second part is the size of the atom. The larger the atom, the more it is able to distribute its electron density which increases the stability of electrons. This takes a precedence over electronegativity, so as we move to the bottom right of the periodic table, acidity will increase. The second factor is resonance, which is one of the most important topics of Organic Chemistry. The more resonance structures a molecule has, the more dispersed the charge is over the molecule making it more stable and acidic. The third factor, and personally the most interesting to me, is induction. When we have an electronegative atom in a molecule, it will pull electrons towards it. This will create an electron poor carbon that is able to pull in other negative charges and stabilize the molecule. This is largely affected by distance, so the closer the electronegative atoms to the negative charge on the molecule, the better. The last factor is orbitals. When there are double or triple bonds on carbon atoms, the hybridization is changed and the total s-character. The higher the s- character percentage the more stable the electrons and the more acidic the atom. These factors are ranked in the order that they were just listed, with one notable exception. NH_2 will be less stable than many other cases. There is one last factor that is discussed in Klein, but was not in our textbook. This factor is the weakest of them all and is based on the shape of the molecules. If all other factors fail, look at the shape of the molecule. If the shape allows for more interactions, meaning it is less bulky, with other molecules, then it will be more acidic. This is the solvating effect. Remember that this is easily understood when you think about the ability for some molecule to reach the hydrogen and react with it. Quantitative Assessment of Acidity We use pKa to assess acidity. The smaller the value, the more acidic a proton is. Remember that a value of 4 versus 7 will be 10^3 times more acidic. 6 vs. 10 is 10^4 more times acidic. Determining the Position of Equilibrium This is a very simple process. You begin with drawing the conjugate bases of the two acids of discussion. Whatever side of the reaction with the more stable conjugate base, will be the side that the equilibrium shifts to. Just remember that everytime you are asked to determine the equilibrium, to draw both conjugate bases and compare. To display an acid-base reaction as well, through the movement of electrons, we use curved arrows like in resonance structures. The only difference is single bonds now can be broken because the molecules are changing. Remember the arrows are dependent on the movement of the electrons, not the movement of the proton.


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