1 Organic Chemistry, Second Edition Introduction to Organic Chemistry Chapter 1 Structure and Bonding Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 2 What is Organic Chemistry? Studies of Organic Compounds Organic Compounds are compounds in (from) living things. Inorganic Compounds are compounds in nonliving things. Historically Until Whler made urea from ammoniumcyanate in organic chemistry is the study of the chemistry of carbon compounds. carbon compounds: natural compounds - terpenoids, alkaloids, carbohydrates, proteins, nucleic acids, vitamines, fats and oils synthetic compounds polymers, plastics organic chemistry is the meeting point of physical science and life science, as it started from life science and became an exact science (physical science). There are more than 20 million organic compounds in the CAS registry today and rapidly increasing! Carbons form 4 strong covalent bonds, to form long chains and rings of many sizes, and to form multiple bonds. Learning organic chemistry is like learning a new language. Unfortunately, we will start with grammar and vocabularies. i.e. it takes time and requires practice to master organic chemistry. Examples of Organic Compounds Organic Chemistry is for . Organic chemistry touches the lives of everyone from the good (for example, medicines and polymers) through the bad (for example, pollution) to the ugly (for example, nerve gases). CH221 Organic Chemistry I Phone : , Office : (E6-4) 4104 lecture notes website :Main textbook : Organic Chemistry, 2nd ed. by Janice Gorzynski Smith Auxiliary textbook : Advanced Organic Chemistry, 4th ed. by F. Carey and R. Sundberg Evaluation Criteria -Final exam : about 30%, midterm exam : about 20%, -Quiz (4 times) : about 40% -Attendance : about 10% Practice Hour : Friday 7:00 9:00 ~ 40minutes of class review ~ 40 minutes of homework review, problem solving, Q&A ( ~ 40 minutes of Quiz ) Lecture Schedule 1st week : Molecular Structure and Bonding (Chap 1) 2nd week : Acids and Bases (Chap 2) 3rd week : Introduction to Organic Molecules and Functional Groups (Chap 3), 1st Quiz 4th week : Alkanes (Chap 4) 5th week : Conformation of Cycloalkanes (Chap 4), Stereochemistry (Chap 5) 6th week : Stereochemistry (Chap 5), Basics of Organic Reactions (Chap 6), 2nd Quiz 7th week : Understanding Organic Reactions (Chap 6) 8th week : Midterm Exam. 9th week : Alkyl Halides (Chap 7), 10th week : Substitution Reaction (Chap 7), Elimination reaction (Chap 8) 11th week : Elimination Reaction (Chap 8), Alcohols (Chap 9), 3rd Quiz 12th week : Alcohols, Ethers (Chap 9) 13th week : Alkenes (Chap 10) 14th week : Alkynes (Chap 11), 4th Quiz 15th week : Oxidation and Reduction (Chap 12) 16th week : Final Exam 8 The nucleus contains positively charged protons and uncharged neutrons. The electron cloud is composed of negatively charged electrons. Structure and Bonding The Periodic Table 9 Elements in the same row are similar in size. Elements in the same column have similar electronic and chemical properties. The Periodic Table Structure and Bonding ElementIsotopes (% natural abundance) Atomic weight (to maximum of 4 decimal places) Hydrogen 1 H (99.98) 2 H (0.02) (= deuterium, D), 3 H (v. small) (= tritium, T) Carbon 12 C (98.9) 13 C (1.1), 14 C (v. small) Nitrogen 14 N (99.63) 15 N (0.37) Oxygen 16 O (99.8) 17 O (0.037) 18 O (0.20) 10 The Periodic Table Structure and Bonding Figure 1.2 A periodic table of the common elements seen in organic chemistry 11 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 of electron density at the nucleus. It is higher in energy than an s orbital. The Periodic Table Structure and Bonding 12 Since there is only one orbital in the first shell, and each shell can hold a maximum of two electrons, there are two elements in the first row, H and He. Each element in the second row of the periodic table has four orbitals available to accept additional electrons: one 2s orbital, and three 2p orbitals. Structure and Bonding 13 Second Row Elements Since each of the four orbitals available in the second shell can hold two electrons, there is a maximum capacity of eight electrons for elements in the second row. The second row of the periodic chart consists of eight elements, obtained by adding electrons to the 2s and three 2p orbitals. Structure and Bonding 14 Review of Bonding Bonding is the joining of two atoms in a stable arrangement. Through bonding, atoms attain a complete outer shell of valence electrons. Through bonding, atoms attain a stable noble gas configuration. Ionic bonds result from the transfer of electrons from one element to another. Covalent bonds result from the sharing of electrons between two nuclei. Structure and Bonding 15 An ionic bond generally occurs when elements on the far left side of the periodic table combine with elements on the far right side, ignoring noble gases. A positively charged cation formed from the element on the left side attracts a negatively charged anion formed from the element on the right side. An example is sodium chloride, NaCl. Structure and Bonding 16 Hydrogen forms one covalent bond. When two hydrogen atoms are joined in a bond, each has a filled valence shell of two electrons. Bonding in Molecular Hydrogen (H 2 ) Structure and Bonding 17 Second row elements can have no more than eight electrons around them. For neutral molecules, this has two consequences: Atoms with one, two, or three valence electrons form one, two, or three bonds, respectively, in neutral molecules. Atoms with four or more valence electrons form enough bonds to give an octet. This results in the following equation: 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. Structure and Bonding 18 Structure and Bonding Figure 1.3 Summary: The usual number of bonds of common neutral atoms 19 Review of Lewis Structures Lewis structures are electron dot representations for molecules. There are three general rules for drawing Lewis structures: In a Lewis structure, a solid line indicates a two-electron covalent bond. 1.Draw only the valence electrons. 2.Give every second-row element an octet of electrons, if possible. 3.Give each hydrogen two electrons. Structure and Bonding 20 Formal Charge cf. Oxidation State Formal charge is the charge assigned to individual atoms in a Lewis structure. By calculating formal charge, we determine how the number of electrons around a particular atom compares to its number of valence electrons. 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. Structure and Bonding 21 The number of electrons owned by different atoms is indicated in the following examples: Example 1 Example 2 Example 3 Structure and Bonding 22 Structure and Bonding 23 Isomers In drawing a Lewis structure for a molecule with several atoms, sometimes more than one arrangement of atoms is possible for a given molecular formula. Example: Isomers are different molecules having the same molecular formula. Ethanol and dimethyl ether are constitutional(structural) isomers (compounds with different connectivity of their atoms). Structure and Bonding 24 Exceptions to the Octet Rule Elements in Groups 2A and 3A Elements in the Third Row Structure and Bonding 25 Resonance structures These look different but are not isomers. Structure and Bonding meaning they are not different compounds 26 Introduction to Resonance Theory 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. Structure and Bonding 1.Resonance structures are not real. An individual resonance structure does not accurately represent the structure of a molecule or ion. Only the hybrid does. 2.Resonance structures are not in equilibrium with each other. There is no movement of electrons from one form to another. 27 Drawing Resonance Structures 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. electron pair = chemical bond or lone pair Example 1: Example 2: Structure and Bonding 28 Resonance Structure Examples In the above examples, an atom bearing a (+) charge is bonded either to a double bond or an atom with a lone pair: Structure and Bonding 29 Resonance Hybrids 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. The hybrid is a weighted average of the contributing resonance structures. Structure and Bonding 30 Drawing Resonance Hybrids v.s. Tautomers Structure and Bonding 31 Determining Molecular Shape Two variables define a molecules structure: bond length and bond angle. Bond length decreases across a row of the periodic table as the size of the atom decreases. Structure and Bonding Bond length increases down a column of the periodic table as the size of an atom increases. 32 Structure and Bonding 33 Bond angle determines the shape around any atom bonded to two other atoms. Determining Molecular ShapeBond Angle 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. Structure and Bonding 34 Determining Molecular ShapeBond Angle Two groups around an atom Structure and Bonding 35 Determining Molecular ShapeBond Angle Three groups around an atom Structure and Bonding 36 Determining Molecular Shape Four groups around an atom Structure and Bonding 37 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. Structure and Bonding 38 Drawing Three Dimensional Structures The molecule can be turned in many different ways, generating many equivalent representations. All of the following are acceptable drawings for CH 4. Structure and Bonding 39 Drawing Three Dimensional Structures Note that wedges and dashes 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, because the two H atoms are really aligned. Structure and Bonding 40 Structure and Bonding Methane (CH 4 ) Ammonia (NH 3 ) Water (H 2 O) In both NH 3 and H 2 O, the bond angle is smaller than the theoretical tetrahedral bond angle because of repulsion of the lone pairs of electrons. The bonded atoms are compressed into a smaller space with a smaller bond angle. 41 Predicting Geometry Based on Counting of Groups Around the Central Atom Structure and Bonding Figure 1.4 Summary: Determining geometry based on the number of groups 42 Drawing Organic MoleculesCondensed Structures 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. Structure and Bonding 43 Examples of Condensed Structures (from Figure 1.5) 44 Examples of Condensed Structures Containing a C-O Double Bond (from Figure 1.6) 45 Skeletal Structures 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 hydrogens directly bonded to them. Structure and Bonding 46 Examples of Skeletal Structures (from Figure 1.7) Structure and Bonding 47 Words of Caution Regarding Interpretation of Skeletal Structures 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. Structure and Bonding 1. Assign Geometry using VSEPR Theory CH 4 SN=4 geometry = tetrahedral 2. Write electronic configuration of the atom to be hybridized C: (1s) 2 (2s) 2 (2p) 2 3. Draw energy diagram from said atom and decouple paired electrons sp 3 Hybrid Make CH 4 Molecular Orbitals in Polyatomic Molecules : Valence Bond Model Valence Bond Model : adopt the idea of Lewis electron pair model. so localize molecular orbitals as electron pairs of bonding or lone pairs. generation of hybrid orbitals from the participating orbitals of atoms. these orbitals overlap to produce bonds ( bond and bond) 49 Shape and Orientation of sp 3 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 (Figure 1.8). The four hybrid orbitals are oriented towards the corners of a tetrahedron, and form four equivalent bonds. Structure and Bonding 50 Bonding Using sp 3 Hybrid Orbitals Each bond in CH 4 is formed by overlap of an sp 3 hybrid orbital of carbon with a 1s orbital of hydrogen. These four bonds point to the corners of a tetrahedron. Structure and Bonding Figure 1.9 Bonding in CH 4 using sp 3 hybrid orbitals 51 Hybridization and Bonding in Organic Molecules 52 Hybridization and Bonding in Organic Molecules Each carbon is trigonal and planar. Each carbon is sp 2 hybridized Structure and Bonding 53 Hybridization and Bonding in Organic Molecules 54 Hybridization and Bonding in Organic Molecules Structure and Bonding 55 Hybridization and Bonding in Organic Molecules Each carbon atom has two unhybridized 2p orbitals that are perpendicular to each other and to the sp hybrid orbitals. 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. 56 Hybridization and Bonding in Organic Molecules Back to resonance structures 57 Bond Length and Bond Strength As the number of electrons between two nuclei increases, bonds become shorter and stronger. Thus, triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds. Structure and Bonding 58 Bond Length and Bond Strength The length and strength of C H bonds vary depending on the hybridization of the carbon atom. Structure and Bonding 59 Structure and Bonding 60 Bond Length and Bond Strength Structure and Bonding 61 Structure and Bonding Note: As the percent s-character increases, a hybrid orbital holds its electrons closer to the nucleus, and the bond becomes shorter and stronger. Although sp 3, sp 2 and sp hybrid orbitals are similar in shape, they are different in size. Bond Length and Bond Strength 62 Electronegativity and Bond Polarity Electronegativity values are used as a guideline 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. When differences in electronegativity result in unequal sharing of electrons, the bond is polar, and is said to have a separation of charge or a dipole. A carboncarbon bond is nonpolar. The same is true whenever two different atoms having similar electronegativities are bonded together. CH bonds are considered to be nonpolar because the electronegativity difference between C and H is small. 63 Electronegativity and Bond Polarity Bonding between atoms of different electronegativity values results in unequal sharing of electrons. Example: In the CO 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, separation of charge. 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, with a perpendicular line drawn through it, is drawn at the less electronegative element. + means the indicated atom is electron deficient. - means the indicated atom is electron rich. 64 Polarity of Molecules Use the following two-step procedure to determine if a molecule has a net dipole: 1.Use electronegativity differences to identify all of the polar bonds and the directions of the bond dipoles. 2.Determine the geometry around individual atoms by counting groups, and decide if individual dipoles cancel or reinforce each other in space. Electrostatic potential plot of CH 3 Cl Structure and Bonding 65 Polarity of Molecules A polar molecule has either one polar bond, or two or more bond dipoles that reinforce each other. An example is water: A nonpolar molecule has either no polar bonds, or two or more bond dipoles that cancel. An example is carbon dioxide: Structure and Bonding 1.11, 1.15, 1.20, 1.28, 1.31, 1.35, 1.42, 1.43, 1.49, 1.56, 1.59, 1.66, 1.70, 1.74, 1.82 Homework