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Organic Chemistry PRINCIPLES AND MECHANISMS Chapter 1: Atomic and Molecular Structure 1.1 What is Organic Chemistry? • Organic chemistry is the branch of chemistry involving organic compounds • In the late 1700s organic compounds were defined as those compounds obtained from a living organism. • Inorganic compounds were defined as those that come from everything else. • The definition for organic chemistry was redefined in 1828 when Friedrich Wöhler synthesized urea (an organic compound from inorganic compounds). Organic Compound Defined • An organic compound is composed primarily of carbon and hydrogen. • This definition is not perfect since some compounds, like CO2, are also considered inorganic. Organic Compound Defined continued… 1.2 Why Carbon? • The carbon atom is capable of having four covalent bonds to other atoms • Carbon atoms can link together in chains of almost any length (branched chains as well as straight chain) • To date, there are tens of millions of organic compounds known an infinite number are possible 1.2 Why Carbon? continued… Silicon Won’t Work • Silicon is found just under carbon on the Periodic Table. • Silicon can also form four covalent bonds. • C–C bond (339 kJ/mol) is stronger than Si-Si bond (188 kJ/mol); less energy is needed to break the Si-Si bond • The strength of typical bonds involving carbon atoms helps biomolecules remain intact—this is essential for molecules to store information or provide cellular structure. • Electronegativity is a better measure of chemistry. Silicon is more metallic (electropositive) than carbon Phosphorus is Better • Phosphorus is actually a closer analog given its electronegativity. • It is a larger and more polarizable element, and therefore oxidizes more readily limiting its ability to act like carbon. • It is an important element in the burgeoning field of organic light emitting diodes (OLEDs) 1.3 Atomic Structure and Ground State Electron Configurations • At the center of an atom there is a positively charged nucleus, composed of protons and neutrons. 1.3 Atomic Structure and Ground State Electron Configurations continued… • Surrounding the nucleus is a cloud of negatively charged electrons, attracted to the nucleus by simple electrostatic forces (the forces by which opposite charges attract one another and like charges repel one another). • The space the electron cloud occupies is far larger than the volume occupied by the nucleus. The Structure of an Atom • An atom has no net charge. • The number of electrons equals the number of protons. • The number of protons in the nucleus is known as the atomic number (Z). • Protons have a +1 charge, electrons have a -1 charge, and neutrons have no charge. • If the number of protons and electrons are unequal, then the entire species bears a charge. – Cation: More protons than electrons. – Anion: More electrons than protons. Atomic Orbitals and Shells • Electrons in an isolated atom reside in atomic orbitals. • The exact location of an electron can never be pinpointed. • The orbital specifies the region of space where the probability of finding a given electron is high. Atomic Orbitals and Shells continued… 1. Atomic orbitals have different shapes. 2. Atomic orbitals are organized in shells which are defined by the principal quantum number, n. 3. Up to two electrons are allowed in any orbital. 4. With increasing shell number, the size and energy of the atomic orbital increases. 5. Within a given shell, an atomic orbital’s energy increases in the following order: s , p , d, etc. In the second shell, for example, the 2s orbital is lower in energy than the 2p. Ground State Electron Configurations: Valence Electrons and Core Electrons • The way in which electrons are arranged in atomic orbitals is called the atom’s electron configuration. • The most stable (i.e., the lowest energy) electron configuration is called the ground state configuration. • Electrons occupying the outermost shell (i.e., highest energy) are valence electrons. The rest are core electrons. The Three Rules to Obtaining an Atom’s Ground State Electron Configuration 1. Pauli’s exclusion principle: No more than two electrons (i.e., zero, one, or two electrons) can occupy a single orbital; two electrons in the same orbital must have opposite spins. 2. Aufbau principle: Each successive electron must fill the lowest-energy orbital available. 3. Hund’s rule: All orbitals at the same energy must contain a single electron before a second electron can be paired in the same orbital. The Ground State Electron Configuration For The First 28 Electrons 1.4 The Covalent Bond: Bond Energy and Bond Length • A covalent bond is characterized by the sharing of valence electrons between two or more atoms, as shown for two H atoms 1.4 The Covalent Bond: Bond Energy and Bond Length continued… • When two H atoms separated by a large distance are brought together, their total energy begins to decrease. • At one particular internuclear distance, the energy of the molecule is at a minimum. 1.4 The Covalent Bond: Bond Energy and Bond Length • Single bonds are the most common type of bond found in organic molecules. • The main difference among single, double, and triple bonds is the number of electrons involved. • In a single bond, two electrons are shared between different nuclei; in a double bond, four electrons are shared; and in a triple bond, six electrons are shared. • As the number of bonds increases between a pair of atoms, bond strength increases and bond length decreases. Average Bond Energies of Common Single Bonds Average Bond Energies and Bond Lengths of Single and Multiple Bonds • Multiple bonds are shorter and stronger than single bonds. Steps to Drawing a Lewis Structure 1. Count the total number of valence electrons in the molecule. a. b. The number of electrons contributed by each atom is the same as its group number (H = 1, C = 4, N = 5, O = 6, F = 7). Each negative charge increases the number of electrons by one; each positive charge decreases the number of electrons by one. 2. Write the skeleton of the molecule, showing only the atoms and the single bonds required to hold them together. a. If molecular connectivity is not given to you, the central atom (the one with the greatest number of bonds) is usually the one with the smallest electronegativity. Steps to Drawing a Lewis Structure continued… 3. Subtract two electrons from the total in Step 1 for each single covalent bond drawn in Step 2. 4. Distribute the remaining electrons as lone pairs. a. b. Start with the outer atoms and work inward. Try to achieve a filled valence shell on each atom—namely, an octet on each atom other than hydrogen, and a duet on each hydrogen atom. 5. If there is an atom with less than a filled valence shell, convert lone pairs from neighboring atoms into bonding pairs of electrons, thereby creating double and/or triple bonds. Covalent Bonding: Sharing Electrons to Produce Full Valence Shells Incomplete Valence Shells and the Expanded Octet • Not all atoms have a complete valence shell. • In borane (BH3), the B atom has a share of only 6 valence electrons (lacks an octet). Incomplete Valence Shells and the Expanded Octet continued… • There are not enough valence electrons available to achieve a complete valence shell for all atoms in the molecule. • Thionyl chloride (SOCl2), the S atom has a share of 10 electrons. 1.6 Strategies for Success: Drawing Lewis Dot Structures Quickly • Drawing Lewis structures can become easier and faster when you observe patterns in the number of bonds for certain atoms. • Each type of atom tends to form a specific number of bonds and to have a specific number of lone pairs of electrons. 1.6 Strategies for Success: Drawing Lewis Dot Structures Quickly continued… 1.7 Electronegativity, Polar Covalent Bonds, and Bond Dipoles • Covalent bonds are characterized by the sharing of electrons between two atomic nuclei. • If the atoms are identical, the electrons are shared equally. • If the atoms are different, than one nucleus will attract electrons more strongly than the other. • That ability to attract electrons in a covalent bond is governed by the element’s electronegativity (EN). • The most well-known EN scale was developed by Linus Pauling, where atoms range from 0 to 4 on electronegativity. Electronegativity Trends • Within the same row, electronegativity values tend to increase from left to right across the periodic table. • Within the same column, they tend to increase from bottom to top. • Elements with the largest electronegativities tend to be in the upper right corner of the periodic table (e.g., N, O, F, Cl, and Br). – These do not include noble gases • Elements with the smallest electronegativities tend to be in the lower left corner (e.g., Rb, Cs, Fr, Ba, and Ra). Electronegativity Trends continued… The Bond Dipole • In a covalent bond, electrons are likely to be found near the nucleus of the more electronegative atom and less likely near the nucleus of the less electronegative atom. • This creates a separation of partial positive and negative charges along the bond, called a bond dipole. The Bond Dipole continued… Electrostatic Potential Maps • Another useful way to illustrate the distribution of charge along a covalent bond is with an electrostatic potential map. • An electrostatic potential map shows a molecule’s electron cloud in colors that indicate its relative charge. Electrostatic Potential Maps continued… • Red corresponds to a buildup of negative charge, whereas blue represents a buildup of positive charge. 1.8 Ionic Bonds • When elements in a compound have large enough differences in electronegativity, ionic bonding can occur. • Rather than sharing electrons, the more electronegative atom acquires electrons given up by the less electronegative atom, forming oppositely charged ions. • The electrostatic attraction between the positively charged cations and the negatively charged anions constitutes the ionic bond. • Sodium chloride (NaCl), for example, consists of sodium cations (Na+) and chloride anions (Cl-). Polyatomic Ions • Polyatomics can include the hydroxide (OH⁻), methoxide (CH3O⁻), and methylammonium (CH3NH3+) ion. • Usually consist only of nonmetals, and their atoms are held together by covalent bonds. 1.9 Assigning Electrons to Atoms in Molecules: Formal Charge and Oxidation State • In an isolated atom or atomic ion, charge is determined by the difference between the atom’s group number and the actual number of valence electrons it possesses. • To get charge, we assign electrons to atoms involved in covalent bonds by two methods: formal charge or oxidation state. • Formal charge: In a given covalent bond, half the electrons are assigned to each atom involved in the bond. • Oxidation state: In a given covalent bond, all electrons are assigned to the more electronegative atom. If the two atoms are identical, the electrons are split evenly. Obtaining Formal Charge and Oxidation State • Formal charges are used much more often than oxidation states. • Unless otherwise stated, you may assume that any charges that appear in a Lewis structure are formal charges. Obtaining Formal Charge and Oxidation State continued… 1.10 Resonance Theory • Some species are not described well by Lewis structures. • In HCO2⁻, both of the C-O bonds are identical experimentally; (i.e., same bond length and bond strength), intermediate between those of a single and a double bond. Resonance Theory Rules • Rule 1 Resonance occurs in species for which there are two or more valid Lewis structures. • Rule 2 Resonance structures are imaginary; the one, true species is represented by the resonance hybrid. Resonance Theory Rules continued… • Rule 3 The resonance hybrid looks most like the lowest energy (most stable) resonance structure. • The two resonance contributors of HCO2⁻ are equivalent. • As a result, each structure contributes equally to the resonance hybrid. • In cases in which resonance structures are inequivalent, their relative energies must be determined in order to obtain their relative contributions to the hybrid. Resonance Theory Rules continued… • A resonance structure is lower in energy (i.e., more stable) with: • a greater number of atoms with filled valence shells • more covalent bonds • fewer atoms with formal charges other than zero • Rule 4 • Resonance provides stabilization. • Resonance structures are stabilized as a result of the delocalization of electrons (electrons have lower energy when they are less confined). • Delocalization is lower in energy than localization of electrons. Delocalization of Electrons Resonance Theory Rules continued… • Rule 5 Resonance stabilization is usually large where resonance structures are equivalent. • If resonance structures are equivalent, then they will contribute equally to the hybrid, allowing electrons the greatest possible delocalization! • Rule 6 All else being equal, the greater the number of resonance structures, the greater the resonance stabilization. Resonance Theory Rules continued… 1.11 Strategies for Success: Drawing All Resonance Structures • It is important to be able to draw all of a species’ resonance structures for two reasons: All resonance structures contribute to the features of the resonance hybrid, and The total number of resonance structures is related to the species’ stability. 1.11 Strategies for Success: Drawing All Resonance Structures continued… • Different resonance structures can be obtained by movement of electrons. • We represent movement of electrons by use of curved arrows. Lone Pair Adjacent to a Multiple Bond • The two curved arrows used to interconvert the resonance structures of HCO2⁻ can be used whenever a Lewis structure exhibits a lone pair of electrons on an atom connected to a double bond or triple bond. • The first curved arrow is used to convert the lone pair into a covalent bond. • The second curved arrow is used to convert a pair of electrons from the original multiple bond into a lone pair. When Formal Charges on an Atom Changes • When a lone pair of electrons on an atom is converted into a bonding pair, the formal charge of that atom becomes more positive by 1. • When a bonding pair involving an atom becomes a lone pair on that atom, the formal charge of that atom becomes more negative by 1. An Incomplete Octet Adjacent to a Multiple Bond • Resonance does not always involve lone pairs. • If an atom with an unfilled valence shell is adjacent to a double bond or triple bond, then we can convert from one resonance structure to another using a single curved arrow. An Incomplete Octet Adjacent to a Multiple Bond • When an atom lacking an octet gains a bond to achieve an octet, the atom’s formal charge becomes more negative by 1. • When a bond is removed from an atom that initially has an octet, the formal charge of the atom becomes more positive by 1. An Incomplete Octet Adjacent to a Multiple Bond continued… A Lone Pair Adjacent to an Atom with an Incomplete Octet • An atom with a lone pair attached to an atom lacking an octet can convert from one resonance structure to another by involving just two atoms. A Lone Pair Adjacent to an Atom with an Incomplete Octet continued… A Ring of Alternating Single and Multiple Bonds • With a ring of alternating single and multiple bonds, a pair of electrons from each multiple bond can shift around the ring—either clockwise or counterclockwise— to arrive at a new resonance structure A Ring of Alternating Single and Multiple Bonds continued… • We must be careful. If only two or four electrons are shifted, we end up with a Lewis structure that is not valid. 1.12 Shorthand Notations • Lone pairs of electrons are frequently omitted. Be careful! • You must have the knowledge of how formal charges relates to the number of bonds and lone pairs! 1.12 Shorthand Notations continued… Shorthand Notations continued… • Condensed formulas allow us to include molecules and molecular ions as part of regular text. • Each nonhydrogen atom is written explicitly, followed immediately by the number of hydrogen atoms that are bonded to it. • Adjacent nonhydrogen atoms in the condensed formula are interpreted as being covalently bonded to each other. Shorthand Notations continued… Shorthand Notations continued… • Line structures, like condensed formulas, are compact and can be drawn quickly and easily. Line Structure Example • The line structure of CH3CH2CH2CH2CH2CH2CH2NH2 is drawn below. 1.13 An Overview of Organic Compounds: Functional Groups • Specific arrangements of atoms connected by specific types of bonds tend to react in characteristic ways. • These structural components are called functional groups. • A molecule’s reactivity is dictated by the functional groups it possesses. • Compounds consisting of only C-C and C-H single bonds are called alkanes. • There are several functional groups, each having their own functional group name. Common Functional Groups Common Functional Groups Biomolecules, Fundamental Building Blocks, and Functional Groups • Organic molecules found exclusively in living organisms are called biomolecules. • There are four major classes of biomolecules: proteins, carbohydrates, nucleic acids, and lipids. Proteins and Amino Acids • Proteins are highly versatile biomolecules. • Actin and myosin are responsible for the mechanical processes involving muscle tissue. • Cortactin is responsible for regulating cell shape. • There are proteins that are classified as enzymes, which act as catalysts for biological reactions. • Proteins are constructed from relatively few types of small organic molecules called a-amino acids. Proteins and Amino Acids continued… Proteins and Amino Acids continued… The 20 Naturally Occurring Amino Acids The 20 Naturally Occurring Amino Acids continued… The 20 Naturally Occurring Amino Acids continued… The 20 Naturally Occurring Amino Acids continued… Carbohydrates and Monosaccharides • Carbohydrates, also called saccharides, serve a variety of biological functions: fuels for metabolic pathways (ex. glycogen) structure components of cell walls (ex. cellulose) involved in blood clotting and can help regulate the immune system • Carbohydrates are characterized by their chemical composition: Carbohydrates and Monosaccharides continued… • Carbohydrates are frequently large molecules, called polysaccharides (which are composed of smaller molecules known as monosaccharides or simple sugars). Carbohydrates and Monosaccharides continued… Nucleic Acids and Nucleotides • A nucleic acid is a large molecular chain that is primarily associated with the storage and transfer of genetic information. • A pair of intertwined nucleic acids form the doublehelical deoxyribonucleic acid (DNA), which stores genetic information, and ribonucleic acid (RNA), which participates in protein synthesis. • Nucleic acids are constructed from relatively small molecular units called nucleotides. Nucleic Acids and Nucleotides continued… Nucleic Acids and Nucleotides continued… All nucleotides have three distinct components: 1. an inorganic phosphate (PO4) group, 2. a cyclic monosaccharide (or sugar), and 3. a nitrogenous base. Nucleic Acids and Nucleotides continued… Nitrogenous Bases in RNA and DNA RNA • The 4 nitrogenous bases of RNA are uracil, guanine, adenine, and cytosine (abbreviated U, G, A, and C, respectively). DNA • The 4 nitrogenous bases of DNA are G, A, C, and thymine (T). • Three of the bases in DNA are the same as the bases in RNA. • Only T (in DNA) is different from U (in RNA). Nitrogenous Bases in RNA and DNA continued… Summary and Conclusions • Organic chemistry has been redefined to focus solely on compounds containing carbon atoms and the reactions these molecules undergo. • Lewis structures, electronegativity, formal charges, and resonance structures help us understand the structure and reactivity of many organic compounds. • Polar covalent bonds arise when atoms with moderate differences in electronegativity are bonded together. Ionic bonds arise when there are large differences in electronegativity. • Shorthand notation allows us to conveniently, accurately, and efficiently draw the structures of simple and complex organic compounds. Summary and Conclusions continued… • Formal charges and oxidation states reflect different ways in which valence electrons are assigned to individual atoms within a molecule. • Resonance exists when two or more valid Lewis structures can be drawn for a given molecular species. • Functional groups are the common bonding arrangement of relatively few atoms, and these groups dictate the behavior of the entire molecule.