Chemistry

Contents Solutions to Problems Chapter 1 Structure and Bonding 1 Chapter 2 Polar Covalent Bonds; Acids and Bases 20 Review Unit 1 38 Chapter 3 Organic Compounds: Alkanes and Their Stereochemistry 41 Chapter 5 An Overview of Organic Reactions 87 Chapter 4 Organic Compounds: Cycloalkanes and Their Stereochemistry 64 Review Unit 2 102 Chapter 6 Alkenes: Structure and Reactivity 106 Chapter 7 Alkenes: Reactions and Synthesis 131 Chapter 8 Alkynes: An Introduction to Organic Synthesis 159 Review Unit 3 182 Chapter 9 Stereochemistry 185 Chapter 10 Organohalides 213 Chapter 11 Reactions of Alkyl Halides: Nucleophilic Substitutions and Eliminations Review Unit 4 262 231 Chapter 12 Structure Determination: Mass Spectrometry and Infrared Spectroscopy 266 Chapter 13 Structure Determination: Nuclear Magnetic Resonance Spectroscору 287 Review Unit 5 314 Chapter 14 Conjugated Dienes and Ultraviolet Spectroscopy 317 Chapter 15 Benzene and Aromaticity 341 Chapter 16 Chemistry of Benzene: Electrophilic Aromatic Substitution 359 Review Unit 6 398 Chapter 17 Alcohols and Phenols 402 Chapter 18 Ethers and Epoxides; Thiols and Sulfides 437 Review Unit 7 465 Chapter 19 Aldehydes and Ketones: Nucleophilic Addition Reactions 468 Chapter 20 Carboxylic Acids and Nitriles 511 Chapter 21 Carboxylic Acid Derivatives: Nucleophilic Acyl Substitution Reactions Review Unit 8 574 Chapter 22 Carbonyl Alpha-Substitution Reactions 578 Chapter 23 Carbonyl Condensation Reactions 607 Chapter 24 Amines and Heterocycles 642 Review Unit 9 684 Chapter 25 Biomolecules: Carbohydrates 687 Chapter 26 Biomolecules: Amino Acids, Peptides, and Proteins 719 Review Unit 10 747 Chapter 27 Biomolecules: Lipids 750 Chapter 28 Biomolecules: Nucleic Acids 776 Chapter 29 The Organic Chemistry of Metabolic Pathways Review Unit 11 817 792 Chapter 30 Orbitals and Organic Chemistry: Pericyclic Reactions Chapter 31 Synthetic Polymers 842 821 Review Unit 12 858 Appendices Functional-Group Synthesis 861 Functional-Group Reactions 866 Reagents in Organic Chemistry 870 Name Reactions in Organic Chemistry 877 Abbreviations 885 Infrared Absorption Frequencies 888 Proton NMR Chemical Shifts 891 Nobel Prize Winners in Chemistry 892 Answers to Review-Unit Questions 901 536 Chapter 1 - Structure and Bonding Chapter Outline I. Atomic Structure (Sections 1.1-1.3). A. Introduction to atomic structure (Section 1.1). B. C. 1. Atoms consist of a dense, positively charged nucleus surrounded by negatively charged electrons. a. The nucleus is made up of positively charged protons and uncharged neutrons. b. The nucleus contains most of the mass of the atom. c. Electrons move about the nucleus at a distance of about 10-m. 2. The atomic number (Z) gives the number of protons in the nucleus. 3. The mass number (A) gives the total number of protons and neutrons. 4. All atoms of a given eiement have the same value of Z. a. Atoms of a given element can have different values of A. b. Atoms of the same element with different values of A are called isotopes. Orbitals (Section 1.2). 1. The distribution of electrons in an atom can be described by a wave equation. a. The solution to a wave equation is an orbital, represented by . b. predicts the volume of space in which an electron is likely to be found. 2. There are four different kinds of orbitals (s., p, d, f). a. The s orbitals are spherical. b. The porbitals are dumbbell-shaped. c. Four of the five d orbitals are cloverleaf-shaped. 3. An atom's electrons are organized into shells. a. The shells differ in the numbers and kinds of orbitals they have. b. Electrons in different orbitals have different energies. c. Each orbital can hold two electrons. 4. The two lowest-energy electrons are in the 1s orbital. a. The 2s orbital is the next in energy. Each p orbital has a region of zero density, called a node. b. The next three orbitals are 2px, 2py and 2pz, which have the same energy. Electron Configuration (Section 1.3). 1. The ground-state electron configuration of an atom is a listing of the orbitals occupied by the electrons of the atom. 2. Rules for predicting the ground-state electron configuration of an atom: a. Orbitals with the lowest energy levels are filled first. The order of filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d. b. Only two electrons can occupy each orbital, and they must be of opposite spin. c. If two or more orbitals have the same energy, one electron occupies each until all are half-full (Hund's rule). Only then does a second electron occupy one of the orbitals. All of the electrons in a half-filled shell have the same spin. II. Chemical Bonding Theory (Sections 1.4 - 1.5). A. Development of chemical bonding theory (Section 1.4). 1. Kekulé and Couper proposed that carbon has four "affinity units" - carbon is tetravalent. 2. Other scientists suggested that carbon can form double bonds, triple bonds and rings. 2 Chapter 1 3. Van't Hoff and Le Bel proposed that the 4 atoms to which carbon forms bonds sit at the corners of a regular tetrahedron. 4. In a drawing of a tetrahedral carbon, a wedged line represents a bond pointing toward the viewer, and a dashed line points behind the plane of the page. B. Covalent bonds. 1. Atoms bond together because the resulting compound is more stable than the individual atoms. 2. 3. a. Atoms tend to achieve the electron configuration of the nearest noble gas. b. Atoms in groups 1A, 2A and 7A either lose electrons or gain electrons to form ionic compounds. c. Atoms in the middle of the periodic table share electrons by forming covalent bonds. The number of covalent bonds formed by an atom depends on the number of electrons it has and on the number it needs to achieve an octet. Covalent bonds can be represented two ways. a. In Lewis structures, bonds are represented as pairs of dots. b. In line-bond structures, bonds are represented as lines drawn between two atoms. 4. Valence electrons not used for bonding are called lone-pair electrons. Lone-pair electrons are represented as dots. C. Valence bond theory (Section 1.5). 1. Covalent bonds are formed by the overlap of two atomic orbitals, each of which contains one electron. The two electrons have opposite spins. 2. Each of the bonded atoms retains its atomic orbitals, but the electron pair of the overlapping orbitals is shared by both atoms. 3. The greater the orbital overlap, the stronger the bond. 4. 5. Bonds formed by the head-on overlap of two atomic orbitals are cylindrically symmetrical and are called o bonds. Bond strength is the measure of the amount of energy needed to break a bond. 6. Bond length is the optimum distance between nuclei. 7. Every bond has a characteristic bond length and bond strength. III. Hybridization (Sections 1.6-1.10). A. sp Orbitals (Sections 1.6, 1.7). B. 1. Structure of methane (Section 1.6). 2. a. When carbon forms 4 bonds with hydrogen, one 2s orbital and three 2p orbitals combine to form four equivalent atomic orbitals (sp³ hybrid orbitals). b. These orbitals are tetrahedrally oriented. c. Because these orbitals are unsymmetrical, they can form stronger bonds than unhybridized orbitals can. d. These bonds have a specific geometry and a bond angle of 109.5°. Structure of ethane (Section 1.7). a. Ethane has the same type of hybridization as oçcurs in methane. b. The C-C bond is formed by overlap of two sp orbitals. c. Bond lengths, strengths and angles are very close to those of methane. sp² Orbitals (Section 1.8). I. If one carbon 2s orbital combines with two carbon 2p orbitals, three hybrid sp² orbitals are formed, and one p orbital remains unchanged. 2. The three sp² orbitals lie in a plane at angles of 120°, and the p orbital is perpendicular to them. 3. Two different types of bonds form between two carbons. a. A o bond forms from the overlap of two sp- orbitals. b. A bond forms by sideways overlap of two p orbitals. c. This combination is known as a carbon-carbon double bond. Structure and Bonding 4. Ethylene is composed of a carbon-carbon double bond and four o bonds formed between the remaining four sp orbitals of carbon and the 1s orbitals of hydrogen. The double bond of ethylene is both shorter and stronger than the C-C bond of ethane. C. sp Orbitals (Section 1.10). 1. If one carbon 2s orbital combines with one carbon 2p orbital, two hybrid sp orbitals are formed, and two p orbitals are unchanged. 2. The two sp orbitals are 180° apart, and the two p orbitals are perpendicular to them and to each other. 3. Two different types of bonds form. a. A o bond forms from the overlap of two sp orbitals. b. Two n bonds form by sideways overlap of four p orbitals. c. This combination is known as a carbon-carbon triple bond. 4. Acetylene is composed of a carbon-carbon triple bond and two o bonds formed between the remaining two sp orbitals of carbon and the 1s orbitals of hydrogen. The triple bond of acetylene is the strongest carbon-carbon bond. D. Hybridization of nitrogen and oxygen (Section 1.10). 1. Covalent bonds between other elements can be described by using hybrid orbitals. 2. Both the nitrogen atom in ammonia and the oxygen atom in water form sp hybrid orbitals. The lone-pair electrons in these compounds occupy sp³ orbitals. 3. The bond angles between hydrogen and the central atom is often less than 109° because the lone-pair electrons take up more room than the o bond. 4. Because of their positions in the third row, phosphorus and sulfur can form more than the typical number of covalent bonds. IV. Molecular orbital theory (Section 1.11). A. Molecular orbitals arise froma mathematical combination of atomic orbitals and belong to the entire molecule. 1. Two 1s orbitals can combine in two different ways. a. The additive combination is a bonding MO and is lower in energy than the two hydrogen 1s atomic orbitals. b. The subtractive combination is an antibonding MO and is higher in energy than the two hydrogen 1s atomic orbitals. 2. A node is a region between nuclei where electrons aren't found. If a node occurs between two nuclei, the nuclei repel each other. 3. The number of MOs in a molecule is the same as the number of atomic orbitals combined. V. Chemical structures (Section 1.12). A. Drawing chemical structures. 1. Condensed structures don't show C-H bonds and don't show the bonds between CH3, CH2 and CH units. 2. Skeletal structures are simpler still. a. Carbon atoms aren't usually shown. b. Hydrogen atoms bonded to carbon aren't usually shown. c. Other atoms are shown. 4 Chapter 1 1.1 Solutions to Problems (a) To find the ground-state electron configuration of an element, first locate its atomic number. For oxygen, the atomic number is 8; oxygen thus has 8 protons and 8 electrons. Next, assign the electrons to the proper energy levels, starting with the lowest level. Fill each level completely before assigning electrons to a higher energy level. Notice that the 2p electrons are in different orbitals. According to Hund's rule, we must place one electron into each orbital of the same energy level until all orbitals are half-filled. 2014 Oxygen 25 4 1s Remember that only two electrons can occupy the same orbital, and that they must be of opposite spin. A different way to represent the ground-state electron configuration is to simply write down the occupied orbitals and to indicate the number of electrons in each orbital. For example, the electron configuration for oxygen is 1s2 2s² 2p4. (b) Silicon, with an atomic number of 14, has 14 electrons. Assigning these to energy levels: 3p 4 + — 35 Silicon 20 H 25# Is# The more concise way to represent ground-state electron configuration for silicon: 152 2s² 2p 3s2 3p² (c) 1s2 2s2 2p6 3s2 3p4 30 4 35 Sulfur 2p H # 2s 4 1s 4 Structure and Bonding 1.2 1.3 Strategy: The elements of the periodic table are organized into groups that are based on the number of outer-shell electrons each element has. For example, an element in group 1A has one outer-shell electron, and an element in group 5A has five outer-shell electrons. To find the number of outer-shell electrons for a given element, use the periodic table to locate its group. Solution: (a) Magnesium (group 2A) has two electrons in its outermost shell. (b) Molybdenum is a transition metal, which has two electrons in the 4s subshell, plus four electrons in its 3d subshell. (c) Selenium (group 6A) has six electrons in its outermost shell. Strategy: A solid line represents a bond lying in the plane of the page, a wedged bond represents abond pointing out of the plane of the page toward the viewer, and a dashed bond represents a bond pointing behind the plane of the page. 1.4 1.5 Solution: H C C CI Chloroform CI H H .H C- H C Ethane H H Strategy: Identify the group of the central element to predict the number of covalent bonds the element can form. Solution: (a) Germanium (Group 4A) has four electrons in its valence shell and forms four bonds to achieve the noble-gas configuration of neon. A likely formula is GeCl4. Element Group Likely Formula (b) Al 3A AlH3 (c) C 4A CH2Cl2 (d) Si 4A SiF4 (e) N 5A CH3NH2 6 Chapter 1 1.6 Strategy: Start by drawing the electron-dot structure of the molecule. (1) Determine the number of valence, or outer-shell electrons for each atom in the molecule. For chloroform, we know that carbon has four valence electrons, hydrogen has one valence electron, and each chlorine has seven valence electrons. 1.7 C 4 x 1 = 4 H 1 x 1 = 1 C 7 x 3 = 21 26 total valence electrons (2) Next, use two electrons for each single bond. H C: C: CI Ci (3) Finally, use the remaining electrons to achieve an noble gas configuration for all atoms.. For a line-bond structure, replace the electron dots between two atoms with a line. Solution: Molecule Electron-dot structure Line-bond structure H H (a) CHCI3 :C:C:CI: :CI-C-CI: :C: :C: (b) H2S H:S: :S-H 1 8 valence electrons H H H H H H (c) CH3NH2 :H:C:N:H H-C-N-H H 14 valence electrons H H H (d) CHgLi H:C:Li 8 valence electrons H H-C-Li H Each of the two carbons has 4 valence electrons. Two electrons are used to form the carbon-carbon bond, and the 6 electrons that remain can form bonds with a maximum of 6 hydrogens. Thus, the formula C2H7 is not possible