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Chapter 1 Structure and Bonding © 2006 Thomson Higher Education Scientific Revolution Scientific revolution leads to: • Safer and more effective medicines • Cures for genetic diseases • Improved quality of life 有机化学（Organic Chemistry ) –研究碳化合物的化学 Carbon Organic chemistry • All organic compounds contain the element of carbon • • • • • Shares four electrons Forms four strong covalent bonds Bonds to other carbons to create chains and rings Not all carbon compounds are derived from living organisms Over 99% of 26 million known compounds contain carbon Chapter 1 Structure and bonding 1.1 Atomic Structure • Atom: (直径2x10-10 m) • Nucleus • Protons (+) 10-14 ~ 10-15 m • neutrons • Electrons (-) 质量忽略不计，在大约 10-10 m 距离的范围内绕 核运动 1. 2 Atomic Structure: Orbitals Quantum Mechanical Model (量子理学模型） Behavior of a specific electron in an atom described by mathematical expression called a wave equation (原子中一个电子的行为可以用波动 方程的数学表达式来描述） Solution of wave equation is called a wave function（波动方程的解称为波函数） Wave function is an orbital（波函数是一个轨道） Orbital denoted by Greek letter psi, A plot of 2 describes volume of space around nucleus the electron is most likely to occupy Electron cloud has no specific boundary. 2 表示原子核周围电子最容易出现的空间区域。 Wave Properties of Electrons • Standing wave vibrates in fixed location. • Wave function, , mathematical description of size, shape, orientation • Amplitude（振幅） may be positive or negative • Node: amplitude is zero _ + - + => Atomic Structure: Orbitals Four different kinds of orbitals for electrons • Denoted s, p, d, and f • s and p orbitals most important in organic chemistry • s orbitals • Spherical, nucleus at center • p orbitals • Dumbbell-shaped, nucleus at middle • d orbitals • Four cloverleaf-shaped and one dumbbell-doughnut Atomic Structure • Atomic number (Z) • Mass number (A) • Atomic weight • Orbital (): s, p, d, & f. px, py, and pz Atomic Structure: Orbitals p Orbital • In each shell, beginning with the second, there are three perpendicular p orbitals, px, py, and pz, of equal energy • Lobes of a p orbital are separated by a region of zero electron density, a node • Each lobe has a different algebraic（代数） sign, + and -, represented by different colors • Algebraic signs are not charges 1.3 Atomic Structure: Electron Configuration（电子构型） Orbitals are grouped in electron shells of increasing size and energy Electron Shell • • • • Each orbital can be occupied by two electrons First shell contains one s orbital, denoted 1s, which holds only two electrons Second shell contains four orbitals, one s orbital (2s) and three p orbitals (2p), which hold a total of eight electrons Third shell contains nine orbitals, one s orbital (3s), three p orbitals (3p), and five d orbitals (3d), which hold a total of 18 electrons Atomic Structure: Electron Configuration P 15 Ground-state electron configuration • The most stable, lowest-energy electron configuration of a molecule or atom Three rules:German 1. Aufbau principle(构造原理） Aufbau meaning "building up • Lowest-energy orbitals fill first: 1s 2s 2p 3s 3p 4s 3d 2. Pauli Exclusion Principle （ 泡利不相容原理） • Electron spin can have only two orientations, up and down • Only two electrons can occupy an orbital, and they must be of opposite spin to have unique wave equations Atomic Structure: Electron Configuration Hund‘s rule （洪特规则） 3. • If two or more empty orbitals of equal energy are available, electrons occupy each orbital with parallel spins until all orbitals have one electron 复习 Three rules:German 1. Aufbau principle Lowest-energy orbitals fill first: 1s 2s 2p 3s 3p 4s 3d 2. Pauli Exclusion Principle Electron spin can have only two orientations, up and down Only two electrons can occupy an orbital, and they must be of opposite spin to have unique wave equations 3. Hund's rule • If two or more empty orbitals of equal energy are available, electrons occupy each orbital with parallel spins until all orbitals have one electron Atomic orbitals Carbon has six electrons and is in row 2 of the periodic table. This means that there are two shells of atomic orbitals available for these electrons. The first shell closest to the nucleus has a single s orbital – the 1s orbital. The second shell has a single s orbital (the 2s orbital) and three p orbitals (3 × 2p). C 1s2 2s2 2px1 2py1 Carbon is in the second row of the periodic table and has six electrons ● aufbau principle ● Pauli exclusion principle ● Hund’s rule. 1s2 2s2 2px1 2py1 复习 Atomic Structure S S s,p S P 1.4 The Nature of Chemical Bonds: Molecular Orbital Theory （化学键的本质：分子轨道理论） σ bonds π bonds Cl2: p-p overlap Constructive overlap along the same axis forms a sigma bond. Cl–Cl => σ bonds σ H–H H2: s-s overlap antibonding MO. bonding MO. H–Cl CH4 Sigma bonds：σ bonds have a circular cross-section and are formed by the head-on overlap of two atomic orbitals. This is a strong interaction and so sigma bonds are strong bonds. A covalent bond binds two atoms together in a molecular structure and is formed when atomic orbitals overlap to produce a molecular orbital – so called because the orbital belongs to the molecule as a whole rather than to one specific atom. antibonding MO. bonding MO. p bonds • • • Additive combination of two 2p orbitals • Leads to formation of a low energy p bonding MO • The p bonding MO is formed by combining p orbital lobes with the same algebraic sign • No node between nuclei Subtractive combination of two 2p orbitals • Leads to formation of a high energy p* antibonding MO • The p* antibonding MO is formed by combining p orbital lobes with different algebraic signs • Node between nuclei Only the bonding p MO is occupied 1.5 The Nature of Chemical Bonds: Valence Bond Theory Valence bond theory • Bonding theory that describes a covalent bond as resulting from the overlap of two atomic orbitals • Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms, thus bonding the two atoms together • • • H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals H-H bond is cylindrically symmetrical Bonds formed by head-on overlap of two atomic orbitals along a line drawn between the nuclei are sigma (s) bonds 1.5 Valence Bond Theory Bond strength • H2 molecule has 436 kJ/mol less energy than the starting 2 H atoms, the product is more stable than the reactant and the H-H bond has a strength of 436 kJ/mol • Conversely, the bond dissociation energy of H2 is 436 kJ/mol because it requires 436 kJ/mol of energy to break the H2 bond 1.5 The Nature of Chemical Bonds: Valence Bond Theory There is an optimum distance between nuclei that leads to maximum stability called the bond length Bond length • The distance between nuclei at the minimum energy point • Because a covalent bond is dynamic, like a spring, the characteristic bond length is the equilibrium distance between the nuclei of two atoms that are bonded to each other 两个原子核之间有一个使键稳 定性最大的距离(74ppm)。 复 习 σ bonds π bonds 1.6 Development of Chemical Theory • In 1858 August Kekulé （凯库勒，德国化学家） and Archibald Couper（库帕， 苏格兰） independently proposed that carbon is tetravalent (always forms four bonds)Emil Erlenmeyer （德国化学家，提出用直短线代表化学键，两划即双键， 三划即三键的表示方法） proposed a carbon-carbon triple bond for acetylene， Alexander Crum Brown （布朗） proposed a carbon-carbon double bond In 1865， Kekulé suggested that carbon chains can double back to form rings of atoms In 1874， Jacobus van’t Hoff荷兰化学家范霍夫 and Joseph Le Bel （勒 贝尔） proposed four atoms to which carbon is bonded sit at the corner of a regular tetrahedron 盲人摸象 1.6 Development of Chemical Theory Atoms bond because the compound that results is more stable and lower in energy than the separate atoms • • Energy is released from the chemical system when a bond forms Energy is consumed by the system when a bond breaks Valence shell • Outer most electron shell of an atom • Eight electrons in valence shell (an electron octet) impart special stability to noble-gas elements in 8A • Main group elements are governed by their tendency to take on electron configuration of the nearest noble gas 1.6 Development of Chemical Theory Ionic compounds • Some elements achieve an octet configuration by gaining or losing electrons • When an electron is gained or lost from a neutral atom an ion is formed • Ions are charged because they have different numbers of protons and electrons • Ions are held together by an electrostatic attraction, like in Na+ Cl-, forming an ionic bond Carbon achieves an octet configuration by sharing electrons Covalent Bond • A bond formed by sharing electrons between atoms Molecule • A neutral collection of atoms held together by covalent bonds 1.6 Development of Chemical Theory Lewis structures (electron-dot structures) • Representations of covalent bonds in molecules Kekulé structures (line-bond structures) 1.6 Development of Chemical Theory Number of covalent bonds depends on how many additional valence electrons needed to reach noble-gas configuration • H (1s) needs one more electron to attain (1s2) • N (2s22p3) needs three more electrons to attain (2s22p6) Lone-pair electrons • Valence-shell electron pairs not used for bonding • Lone-pair electrons can act as nucleophiles and react with electrophiles 1.7 Hybridization a carbon atom has two unpaired electrons and so we would expect carbon to form two bonds.However, carbon forms four bonds! when a carbon atom forms bonds,it can ‘mix’ the s and p orbitals of its second shell (the valence shell). This is known as hybridization and it allows carbon to form the four bonds which we observe in reality.There are three ways in which this mixing process can take place. ● the 2s orbital is mixed with all three 2p orbitals. This is known as sp3 hybridization; ● the 2s orbital is mixed with two of the 2p orbitals. This is known as sp2 hybridization; ● the 2s orbital is mixed with one of the 2p orbitals. This is known as sp hybridization. sp3 hybridization Energy levels M Methane molecule forms 4 C−H sigma bonds and e t is three dimensional tetrahedral shape，bond h angleais 109.5°，cabon atom is sp3 hybridized in n methane molecule . e sp3 Hybrid Orbitals and the Structure of Ethane Orbital hybridization accounts for the bonding together of carbon atoms into chains and rings Ethane C2H6 • Tetrahedral • Bond angles are near 109.5º • Carbon-carbon single bond • Formed by s overlap of sp3 hybrids from each carbon • The remaining sp3 hybrids of each carbon overlap with 1s orbitals of three hydrogen atoms to form six carbon-hydrogen bonds Ethane: CH3-CH3 Question： σ bonds C-H C-C? 3 Sigma bonds and one lone pair :NH3 Geometry? Pyramidal,why? 2 sigma bonds and 2 lone pair O H H BENT sp2 Hybrid Orbitals 有那些键？ sp2 Hybrid Orbitals and the Structure of Ethylene Ethylene C2H4 • • • Carbon-carbon double bond • Four shared electrons Planar (flat) Bond angles 120º sp2 hybrid orbitals • A hybrid orbital derived by combination of an s atomic orbital with 2p atomic orbitals • One p orbital remains nonhybridized πBonding • Pi bonds form after sigma bonds. • Sideways overlap of parallel p orbitals. => Each ethene molecule forms 1 C–C σ bond , 4 C–H σ bond and 1 π bond. Ethene is a flat, rigid molecule where each carbon is trigonal planar. Carbon atom is sp2 hybridized and bond angle is 120°in ethene molecule. • Ethylene • • • C2H4 s bond π bond Multiple Bonds • A double bond (2 pairs of shared electrons) consists of a sigma bond and a pi bond. • A triple bond (3 pairs of shared electrons) consists of a sigma bond and two pi bonds. sp2 Hybrid Orbitals sp2 Hybrid Orbitals • combination of an s orbital with 2 pi orbitals - three sp2 hybrid orbitals trigonal planar: 120o bond angles. O H 3C H Carbon has 3 sigma and one pi bond O Delocalization occurs in 1,3-butadiene where there are alternating single and double bonds. All four carbons in 1,3-butadiene are sp2 hybridized and so each of these carbons has a half-filled p orbital which can interact to give two π bonds. However, a certain amount of overlap is also possible between the p orbitals of the middle two carbon atoms and so the bond connecting the two alkenes has some double bond character borne out by the observation that this bond is shorter in length than a typical single bond. This delocalization also results in increased stability. SP HYBRIDIZATION • Acetylene • • • C2H2 One s-bond & two p-bonds => Each ethyne molecule forms 1 C–C σ bond , 2 C–H σ bond. The 2py and 2pz orbitals of each carbon atom can overlap side-on to form two π bonds. Each carbon is sp hybridized and bond angle is 180°in ethyne. Ethyne is linear. sp Hybrid Orbitals and the Structure of Acetylene Acetylene • • Linear Carbon-carbon triple bond • Six shared electrons • Bond angles are 180º sp hybridized orbital • A hybrid orbital derived from the combination of one s and one p atomic orbital • The two sp hybrids are oriented at an angle of 180º to each other • Two 2p orbitals remain non-hybridized HCN sp Hybrid Orbitals Combine one 2 s with one 2 p sigma orbital Linear 180 o H-C N Note: both C and N are sp hybridized H-C N sp Hybrid Orbitals and the Structure of Acetylene 复习 ：说出下列分子以及杂化类型 22_502 H H H H C H C H 22_504 2p 2p 2p H1s H1s 2p H C H C sp 2p 2p 2p sp 2p sp 22_501 sp2 sp2 sp2 H1s sp 2 sp2 C C sp2 H1s 2p A double bond is a s bond and a p bond. Double bond B.D.E. = 146 kcal/mol s bond B.D.E. = 83 kcal/mol Therefore p B.D.E. must = 63 kcal/mol. A p bond is weaker than a s bond. Bonds are more reactive than s bonds, p bonds are considered to be a functional group. Double and Triple Bonds H 1.33A H H H H HH 1.08A 1.09A HH H 1.54A sp2 are 1/3 s, whereas sp3 are 1/4 s in character. (s orbitals are closer to the nucleus and lower in energy).The carbon-carbon bond in ethene is shorter and stronger than in ethane partly because of the sp2-sp2 overlap being stronger than sp3-sp3, but especially because of the extra p bond in ethene. σ and π bonds Identifying σ and π bonds in a molecule is quite easy as long as you remember the following rules: ● all bonds in organic structures are either sigma (σ) or pi (π) bonds; ● all single bonds are σ bonds; ● all double bonds are made up of one σ bond and one π bond; ● all triple bonds are made up of one σ bond and two π bonds Hybridized centers: All the atoms in an organic structure (except hydrogen) are either sp, sp2 or sp3 hybridized 含双键和三键的分子比含有单键的分子活泼得多。 1.8 Polar Covalent Bonds: Electronegativity • Covalent bonds can have ionic character • These are polar covalent bonds • Bonding electrons attracted more strongly by one atom than by the other • Electron distribution between atoms is not symmetrical 72 Bond Polarity and Electronegativity • Electronegativity (EN): intrinsic ability of an atom to • • • • • • attract the shared electrons in a covalent bond Differences in EN produce bond polarity Arbitrary scale. As shown in Figure 2.2, electronegativities are based on an arbitrary scale F is most electronegative (EN = 4.0), Cs is least (EN = 0.7) Metals on left side of periodic table attract electrons weakly, lower EN Halogens and other reactive nonmetals on right side of periodic table attract electrons strongly, higher electronegativities EN of C = 2.5 73 The Periodic Table and Electronegativity 74 Bond Polarity and Inductive Effect （键的极化和诱导效应） • • • • • 诱导效应（Inductive Effect） ：一个原子使键极化的能力。 Nonpolar Covalent Bonds: atoms with similar EN Polar Covalent Bonds: Difference in EN of atoms < 2 Ionic Bonds: Difference in EN > 2 • C–H bonds, relatively nonpolar C-O, C-X bonds (more electronegative elements) are polar Bonding electrons toward electronegative atom • C acquires partial positive charge, + • Electronegative atom acquires partial negative charge, Inductive effect: shifting of electrons in a bond in response to EN of nearby atoms 75 Electrostatic Potential Maps 分子的静电势图 • Electrostatic potential maps show calculated charge distributions • Colors indicate electronrich (red) and electronpoor (blue) regions • Arrows indicate direction of bond polarity 体系中静电势大的部位容易受亲电试剂 的攻击发生反应。 76 Bond Polarity & Electronegativity • Covalent bond to ionic bond EN = 0 0< EN <2.0 EN>2.0 1.8 Bond Polarity & Electronegativity • Electrostatic potential maps • Electron rich (red) • Electron poor (blue) 1.9 Resonance • • • • Some molecules are have structures that cannot be shown with a single representation In these cases we draw structures that contribute to the final structure but which differ in the position of the p bond(s) or lone pair(s) Such a structure is delocalized and is represented by resonance forms The resonance forms are connected by a double-headed arrow 1.9 Resonance Hybrids • • • A structure with resonance forms does not alternate between the forms Instead, it is a hybrid of the two resonance forms, so the structure is called a resonance hybrid For example, benzene (C6H6) has two resonance forms with alternating double and single bonds • In the resonance hybrid, the actual structure, all its C-C bonds are equivalent, midway between double and single 80 1.9 Rules for Resonance Forms • Individual resonance forms are imaginary - the real • • • • structure is a hybrid (only by knowing the contributors can you visualize the actual structure) Resonance forms differ only in the placement of their p or nonbonding electrons Different resonance forms of a substance don’t have to be equivalent Resonance forms must be valid Lewis structures: the octet rule applies The resonance hybrid is more stable than any individual resonance form would be 81 1.9 Curved Arrows and Resonance Forms • We can imagine that electrons move in pairs to convert from one resonance form to another • A curved arrow shows that a pair of electrons moves from the atom or bond at the tail of the arrow to the atom or bond at the head of the arrow 82 1.9 Drawing Resonance Forms • Any three-atom grouping with a multiple bond has two resonance forms 83 1.9 Different Atoms in Resonance Forms • • • • Sometimes resonance forms involve different atom types as well as locations The resulting resonance hybrid has properties associated with both types of contributors The types may contribute unequally The “enolate” derived from acetone is a good illustration, with delocalization between carbon and oxygen 84 2,4-Pentanedione • The anion derived from 2,4-pentanedione • Lone pair of electrons and a formal negative charge on the central carbon atom, next to a C=O bond on the left and on the right • Three resonance structures result 85 86 87 Summary • • • • Organic molecules often have polar covalent bonds as a result of unsymmetrical electron sharing caused by differences in the electronegativity of atoms The polarity of a molecule is measured by its dipole moment, . (+) and () indicate formal charges on atoms in molecules to keep track of valence electrons around an atom Some substances must be shown as a resonance hybrid of two or more resonance forms that differ by the location of electrons.