Organic Chemistry Chapter 1 PDF

Overall course discussion CHM 221; Dr. Michael Lucarelli Attendance Syllabus: Policies, exams, Homework Class lecture: Thursdays: ~ 3 hrs Presentations will follow the textbook: Note: “Check my speed” 1 General aspects of course Concepts of scie...

Overall course discussion CHM 221; Dr. Michael Lucarelli Attendance Syllabus: Policies, exams, Homework Class lecture: Thursdays: ~ 3 hrs Presentations will follow the textbook: Note: “Check my speed” 1 General aspects of course Concepts of science and organic chemistry Relate the chemistry to various industries and academia Understand the Scientific Method Homework Participation in discussions: 2  Organic Chemistry Study of the structure, properties, composition, reactions and preparations of carbon-containing compounds, which include hydrocarbons and compounds that contain any number of other elements including hydrogen, nitrogen, oxygens, halogens, phosphorous, silicon and sulfur (and others. Most compounds contain a hydrogen-carbon bond… (American Chemical Society) 3  Concerns Area of “biggest” problems: 1. Nomenclature 2. Reactions: a. Don’t memorize each “molecule” b. “learn” to recognize the reactions and functional groups (next slide as example) c. Don’t let chemistry reactions confuse you….don’t “overthink” questions 3. Stereochemistry: Spatial arrangement of atoms 4  Reaction example Similar “dehydration” reaction regardless of molecule Both lose water (H2O) 5 Don’t let reactions confuse you Murov, S. J. Chem. Ed., 2007, 84(2), 1224 Notice referencing 6  Organic Chemistry Lectures will correspond to textbook Chapter 1 Review of atoms, molecules, bonding, etc Chapter 2 will not be covered: some basics concepts from it at the end of this lecture 7  Atomic Structure 1 The nucleus contains positively charged protons and uncharged neutrons. The electron cloud is composed of negatively charged electrons. 8  Atomic Structure 2 atomic number = number of protons mass number = number of protons plus neutrons In a neutral atom, the number of protons equals the number of electrons. The atomic weight of a particular element is the weighted average of the mass of all its isotopes reported in atomic mass unit (amu). 9  Ions In addition to neutral atoms, we will also encounter charged ions. A cation is positively charged and has fewer electrons than protons. An anion is negatively charged and has more electrons than protons. 10  Isotopes Isotopes are two atoms of the same element having a different number of neutrons. Isotopes have different mass numbers. 11  The Periodic Table Elements in the same row are similar in size. Elements in the same column have similar electronic and chemical properties. Figure 1.1 12  Atomic Orbitals An s orbital has a sphere of electron density and is lower in energy than the other orbitals of the same shell. A p orbital has a dumbbell shape and contains a node (no electron density) at the nucleus. It is higher in energy than an s orbital. 13  Periodic Table 1 The First Row There is only one orbital in the first shell. Each shell can hold a maximum of two electrons. Therefore, there are two elements in the first row: H and He. 14  Periodic Table 2 The Second Row Each element in the second row of the periodic table also has four orbitals available to accept additional electrons: one 2s orbital, and three 2p orbitals. 15  Periodic Table 3 The Second Row Each of the four orbitals in the second shell hold two electrons. There is a maximum capacity of eight valence electrons for elements in the second row. The second row of the periodic table consists of eight elements, obtained by adding electrons to the 2s and three 2p orbitals. 16  Bonding Bonding is the joining of two atoms in a stable arrangement. Through bonding, atoms gain, lose, or share electrons to attain the electronic configuration of the noble gas closest to them in the periodic table. Atoms can form either ionic or covalent bonds to attain a complete outer shell (octet rule for second row elements). Ionic bonds result from the transfer of electrons from one element to another. Covalent bonds result from the sharing of electrons between two nuclei. 17  Covalent Bonding 1 Covalent bonding occurs with elements that have similar electronegativity (non- metals). Covalent bonds also occur between two of the same elements. A covalent bond contains 2 electrons A compound with covalent bonds is called a molecule. 18  Covalent Bonding 2 Bonding in Molecular Hydrogen (H2) Hydrogen forms one covalent bond. When two hydrogen atoms are joined in a bond, each has a filled valence shell of two electrons. 19  Valence Electrons Second row elements can have no more than eight electrons around them. For neutral molecules, this has two consequences: Atoms with one, two, three, or four valence electrons form one, two, three, or four bonds, respectively, in neutral molecules (e.g., BF3, CH4). Atoms with five or more valence electrons form enough bonds to give an octet (e.g., NH3). In this case, predicted number of bonds = 8 − number of valence electrons. 20  Nonbonded Electrons When second-row elements form fewer than four bonds their octets consist of both bonding (shared) and nonbonding (unshared) electrons. Unshared electrons are also called lone pairs. Figure 1.2 21  Lewis Structures (use of Periodic Table) Lewis structures are electron dot representations for molecules. We won’t spend time on understanding how to draw Lewis structures (but you should learn how to draw them for future reference; general chem) In a Lewis Structure, a solid line indicates a two-electron covalent bond. Organic Chemistry generally deals with the covalent bonding of carbon atoms to other carbon atoms and hydrogens, nitrogen, oxygen, halides… Carbon (and other atoms) can make single and multiple bonds 22  Formal Charge dealing with polarity Formal charge is the charge assigned to individual atoms in a Lewis structure. Formal charge is calculated as follows: The number of electrons “owned” by an atom is determined by its number of bonds and lone pairs. An atom “owns” all of its unshared electrons and half of its shared electrons. 23  Electron Ownership 1 The number of electrons “owned” by different atoms is indicated in the following examples: 24  Isomers Sometimes more than one arrangement of atoms (Lewis structure) is possible for a given molecular formula. These two compounds are called isomers. Isomers are different molecules having the same molecular formula. Ethanol and dimethyl ether are constitutional isomers. This is common in organic chem 25  Exceptions to the Octet Rule Elements in Groups 2A and 3A Elements in the Third Row involved with carbon 26  Resonance Some molecules cannot be adequately represented by a single Lewis structure. For example: These structures are called resonance structures or resonance forms. A double-headed arrow is used to separate the two resonance structures. Resonance structures are two Lewis structures having the same placement of atoms but a different arrangement of electrons. 27  Resonance Forms Neither resonance structure is an accurate representation for (HCONH)−. The true structure is a composite of both resonance forms and is called a resonance hybrid. The hybrid shows characteristics of both structures. Resonance allows certain electron pairs to be delocalized over two or more atoms, and this delocalization adds stability. A molecule with two or more resonance forms is said to be resonance stabilized. 28  Basic Principles of Resonance Theory Resonance structures are not real. An individual resonance structure does not accurately represent the structure of a molecule or ion. Only the hybrid does. Resonance structures are not in equilibrium with each other. There is no movement of electrons from one form to another. Resonance structures are not isomers. Two isomers differ in the arrangement of both atoms and electrons, whereas resonance structures differ only in the arrangement of electrons. 29  Drawing Resonance Structures 1 Rule : Two resonance structures differ in the position of multiple bonds and nonbonded electrons. The placement of atoms and single bonds always stays the same. 30  Drawing Resonance Structures 2 Rule : Two resonance structures must have the same number of unpaired electrons. Rule : Resonance structures must be valid Lewis structures. Hydrogen must have two electrons and no second-row element can have more than eight electrons. 31  Curved Arrow Notation pushing electrons for reactions…mechanisms reactions occur if electrons can move…. Curved arrow notation is a convention that shows how electron position differs between two resonance forms. Curved arrow notation shows the movement of an electron pair. The tail of the arrow always begins at the electron pair, either in a bond or lone pair. The head points to where the electron pair “moves.” Example 1: Example 2: 32  Atoms Without Octets we will discuss later in more detail Resonance structures can have an atom with fewer than 8 electrons. However, resonance structures can never have a second- row element with more than 8 electrons. 33  Occurrence of Resonance Two different resonance structures can be drawn when a lone pair is located on an atom directly bonded to a double bond. 34  The Resonance Hybrid A resonance hybrid is a composite of all possible resonance structures. In the resonance hybrid, the electron pairs drawn in different locations in individual resonance forms are delocalized. When two resonance structures are different, the hybrid looks more like the “better” resonance structure. The “better” resonance structure is called the major contributor to the hybrid, and all others are minor contributors. 35  Resonance Hybrids A “better” resonance structure is one that has more bonds and fewer charges. 36 Determining Molecular Shape generalize shape and geometry 1 Two variables define a molecule’s structure: bond length and bond angle. Bond length decreases across a row of the periodic table as the size of the atom decreases. Bond length increases down a column of the periodic table as the size of an atom increases. 37  Determining Molecular Shape 2 Table 1.2 Average Bond Lengths Bond Length(pm) Bond Length(pm) Bond Length(pm) H-H 74 H-F 92 C-F 133 C-H 109 H-Cl 127 C-Cl 177 N-H 101 H-Br 141 C-Br 194 O=H 96 H-I 161 C-I 213 38  Molecular Geometry The number of groups surrounding a particular atom determines its geometry. A group is either an atom or a lone pair of electrons. The most stable arrangement keeps these groups as far away from each other as possible. This is exemplified by Valence Shell Electron Pair Repulsion (VSEPR) theory. Single bonds (sigma bonds) determine geometry: a lone pair acts as a single bond Number of groups Geometry Bond angle two groups linear 180o (2 single bonds) three groups trigonal planar 1200 (3 single bonds) four groups tetrahedral 109.5o (4 single bonds) 39 Two Groups Around an Atom Note: Ball and stick models are emphasized in book as these are used in labs….Concentrate on the drawn models as we go through the lectures…… 40 Three Groups Around an Atom 41 Four Groups Around an Atom 42  Drawing Three-Dimensional Structures A solid line is used for a bond in the plane. A wedge is used for a bond in front of the plane. A dashed line is used for a bond behind the plane 3-D structure is important for understanding/explaining how molecules may react……… 43  Equivalent Representations for Methane 3d orientation The molecule can be turned in many different ways, generating equivalent representations. All of the following are acceptable drawings for CH4. Each drawing has two solid lines, one wedge, and one dashed wedge. 44  Wedges and Dashed Wedges Note that wedges and dashed wedges are used for groups that are really aligned one behind another. It does not matter in the following two drawings whether the wedge or dash is skewed to the left or right. 45  A Nonbonded Pair of Electrons is Counted as a “Group” In ammonia (NH3), one of the four groups attached to the central N atom is a lone pair. The group geometry is a tetrahedron. The molecular shape is referred to as trigonal pyramidal. 46  The 3-D Structure of Water In water (H2O), two of the four groups attached to the central O atom are lone pairs. The group geometry is a tetrahedron. The molecular shape is referred to as bent. 47  Varying Bond Angles In both NH3 and H2O the Methane (CH4) bond angle is smaller than the theoretical tetrahedral bond angle because of repulsion of the lone pairs of Ammonia (NH3) electrons. The bonded atoms are compressed into a smaller space with a Water (H2O) smaller bond angle. ©2020 McGraw-Hill Education. 48  Summary: Predicting Geometry Based on Number of Groups Summary: Determining Geometry Based on the Number of Groups Single Bonds Number of groups around the Lone pair act Around an atom Geometry atom as single bond 2 linear 2 3 trigonal planar 3 4 tetrahedral 4 49 Language of Organic Chemistry: How chemists write and talk about organic molecules………..eventually reflects on chemical reactions 1. Condensed structure 2. Skeletal Structure 3. Hybridization 50  Drawing Organic Molecules—Condensed Structures important All atoms are drawn in, but the two-electron bond lines are generally omitted. Atoms are usually drawn next to the atoms to which they are bonded. Parentheses are used around similar groups bonded to the same atom. Lone pairs are omitted. 51  Examples of Condensed Structures Figure 1.3 52  Condensed Structures with C=O Figure 1.4 In these examples, the only way for all atoms to have an octet is by having a carbon-oxygen double bond. 53  Skeletal Structures 1 Assume there is a carbon atom at the junction of any two lines or at the end of any line. Assume there are enough hydrogens around each carbon to make it tetravalent. Draw in all heteroatoms and the hydrogens directly bonded to them. 54 Skeletal Structures 2 55  Interpreting Skeletal Structures Figure 1.5 56  Skeletal Structures with Charged Carbon Atoms A charge on a carbon atom takes the place of one hydrogen atom. The charge determines the number of lone pairs. Negatively charged carbon atoms have one lone pair and positively charged carbon atoms have none. 57  Lone Pairs on Heteroatoms Skeletal structures often leave out lone pairs on heteroatoms, but don’t forget about them. Use the formal charge to determine the number of lone pairs. 58  Orbitals and Bonding: Hydrogen When the 1s orbital of one H atom overlaps with the 1s orbital of another H atom, a sigma (σ) bond that concentrates electron density between the two nuclei is formed. All single bonds are  bonds. This bond is cylindrically symmetrical because the electrons forming the bond are distributed symmetrically about an imaginary line connecting the two nuclei. 59  Orbitals and Bonding: Methane To account for the bonding patterns observed in more complex molecules, we must take a closer look at how the 2s and 2p orbitals of atoms in the second row are utilized. In addition to its two core electrons, carbon has four valence electrons. In its ground state, carbon places two electrons in the 2s orbital and one each in 2p orbitals. Note: The lowest energy arrangement of electrons for an atom is called its ground state. 60  Hybrid Orbitals To solve this dilemma, chemists have proposed that atoms like carbon do not use pure s and pure p orbitals in forming bonds. Instead, atoms use a set of new orbitals called hybrid orbitals. Hybridization is the combination of two or more atomic orbitals to form the same number of hybrid orbitals, each having the same shape and energy. 61  Shape and Orientation of sp3 Hybrid Orbitals The mixing of a spherical 2s orbital and three dumbbell shaped 2p orbitals together produces four hybrid orbitals, each having one large lobe and one small lobe. The four hybrid orbitals are oriented towards the corners of a tetrahedron, and form four equivalent bonds. 62  Bonding Using sp3 Hybrid Orbitals Each bond in CH4 is formed by overlap of an sp3 hybrid orbital of carbon with a 1s orbital of hydrogen. These four bonds point to the corners of a tetrahedron. Figure 1.6 63  Other Hybridization Patterns Figure 1.7 Figure 1.8 ©2020 McGraw-Hill Education. 64  Determining Hybridization Count the number of groups (atoms and nonbonded electron pairs) around the atom. The number of groups corresponds to the number of atomic orbitals that must be hybridized to form the hybrid orbitals. Table 1.4 Three types of Hybrid Orbitals Number of groups Number of orbitals Type of hybrid orbital 2 2 two sp hybrid orbitals 3 3 three 4 4 four 65  Hybrid orbitals of NH3 and H2O Figure 1.9 66 Hybridization and Bonding in Ethane 67  Energy levels and orbitals s orbital (group 1A, 2A) p orbitals (groups 3A – 8A) n 1 2 d orbital block 3 4 5 6 f orbitals 68 6  Ethane, CH3—CH3 Making a model of ethane illustrates one additional feature about its structure. Rotation occurs around the central C—C  bond. 69 Hybrid Orbitals in Ethylene Each carbon is trigonal planar. Each carbon is sp2 hybridized. 70  σ and π Bonds in Ethylene Figure 1.10 71  No Free Rotation in Ethylene Unlike the C—C bond in ethane, rotation about the C—C double bond in ethylene is restricted. It can only occur if the π bond first breaks and then reforms, a process that requires considerable energy. 72 sp Hybrid Orbitals 73  Acetylene (Ethyne) Each carbon atom has two unhybridized 2p orbitals that are perpendicular to each other and to the sp hybrid orbitals. 74  Triple Bonds The side-by-side overlap of two 2p orbitals on one carbon with two 2p orbitals on the other carbon creates the second and third bonds of the triple bond. All triple bonds are composed of one sigma and two pi bonds. Figure 1.11 75  Summary of Bonding in Acetylene 1 Figure 1.11 76 Summary of Bonding in Acetylene 2 77  Bond Length and Bond Strength As the number of electrons between two nuclei increases, bonds become shorter and stronger. Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds. 78  Carbon-Hydrogen Bonds 1 The length and strength of C—H bonds vary depending on the hybridization of the carbon atom. 79 Carbon-Hydrogen Bonds 2 80 Percent s-Character 81  Electronegativity Electronegativity is a measure of an atom’s attraction for electrons in a bond. Electronegativity values for some common elements: Figure 1.12 82  Bond Polarity Electronegativity values are used to indicate whether the electrons in a bond are equally shared or unequally shared between two atoms. When electrons are equally shared, the bond is nonpolar. 83  Nonpolar Bonds A carbon—carbon bond is nonpolar. C—H bonds are considered to be nonpolar because the electronegativity difference between C and H is small. Whenever two different atoms having similar electronegativities are bonded together, the bond is nonpolar. 84  Polar Bonds Bonding between atoms of different electronegativity values results in unequal sharing of electrons. Example: In the C—O bond, the electrons are pulled away from C (2.5) toward O (3.4), the element of higher electronegativity. The bond is polar, or polar covalent. The bond is said to have dipole; that is, partial separation of charge. A C—O bond is a polar bond. 85  Depicting Polarity The δ+ means the indicated atom is electron deficient. The δ− means the indicated atom is electron rich. The direction of polarity in a bond is indicated by an arrow with the head of the arrow pointing towards the more electronegative element. The tail of the arrow is drawn at the less electronegative element. 86  Polarity of Molecules Use the following procedure to determine if a molecule has a net dipole: Use electronegativity differences to identify all of the polar bonds and the directions of the bond dipoles. Determine the geometry around individual atoms by counting groups, and decide if individual dipoles cancel or reinforce each other in space. Figure 1.13 Electrostatic potential plot of CH3Cl 87  Polar Molecules A polar molecule has either one polar bond, or two or more bond dipoles that reinforce each other. An example is water: It is a bent molecule Two dipoles reinforce It has a net dipole, making it a polar molecule 88  Nonpolar Molecules A nonpolar molecule has either no polar bonds, or two or more bond dipoles that cancel. An example is carbon dioxide: It is a linear molecule Two dipoles are equal and opposite in direction Two dipoles cancel It is a nonpolar molecule with no net dipole 89  Electrostatic Potential for Polar and Nonpolar Molecules The dipoles of H2O and CO2 can also be visualized using electrostatic potential plots. Figure 1.14 90  Polarity can “indicate” acidity Chapt 2 covers acidity…we will not go into any detail, but discuss acidity as we go through reactions… Note next 2 slides….inductive effects 91  Inductive Effects An inductive effect is the pull of electron density through σ bonds caused by electronegativity differences of atoms. More electronegative atoms stabilize regions of high electron density by an electron withdrawing inductive effect. The more electronegative the atom and the closer it is to the site of the negative charge, the greater the effect. The acidity of H-A increases with the presence of electron withdrawing groups in A. 92  Inductive Effects in Trifluoroethanol In the example below, note that 2,2,2-trifluoroethanol is more acidic than ethanol. This is because the three electronegative fluorine atoms stabilize the negatively charged conjugate base. 93 End HW: 58; 59a, d;61;64a,d;67;71 Look over #62, 63, 64 to make sure you understand what the structures represent 94