附件2 粒大浮 教 案 2003~~2004学年第Ⅰ学期 院(系、所、部)化学与环境学院有机化学研究所 教研室有机化学 课程名称有机化学(双语教学 授课对象化学教育 授课教师杨定乔 职称职务教授 教材名称 Organic Chemistry 2003年09月01日
附件 2 教 案 2003~~ 2004 学年 第 I 学期 院（系、所、部）化学与环境学院有机化学研究所 教 研 室 有机化学 课 程 名 称 有机化学(双语教学) 授 课 对 象 化学教育 授 课 教 师 杨定乔 职 称 职 务 教授 教 材 名 称 Organic Chemistry 2003 年 09 月 01 日
有机化学(双语教学)课程教案 授课题目(教学章节或主题):第一章。绪论授课类型理论课 Introduction 授课时间第1周第12节 教学目标或要求:了解基本有机价键理论以及杂化轨道理论 教学内容(包括基本内容、重点、难点) Atomic molecular orbitals Electrons surrounding atoms are concentrated into regions of space called atomic orbitals. The Heisenberg uncertainty principle states that it impossible to know both the location and the momentum of an atomic particle, but it is possible to describe the probability that the electron will be found within a given region of space. The boundaries of an atomic orbital are commonl drawn to the region of 90% probability: there is a 90% probability that at any given time, the electron will be within the specified boundary. The electronic configuration of carbon is ls 2s 2sp. Atomic orbitals with s-character have spherical symmetry and a representation of the surface of the carbon is orbital is shown below 1s atomic orbital The wave properties of electrons make the description of the 2s orbital slightly more complex than the corresponding ls orbital, in that, within the 2s sphere there is a region in which the amplitude of the electron standing wave falls to zero, that is, there is zero probability of finding the electron in this node region. Nodes are most easily seen in the description of the 2p atomic orbitals, which are shown below
有机化学（双语教学） 课程教案 授 课 题 目（ 教 学 章节 或 主题 ）：第 一 章 。绪 论 （Introduction） 授课类型 理论课 授课时间 第 1 周第 1-2 节 教学目标或要求：了解基本有机价键理论以及杂化轨道理论。 教学内容（包括基本内容、重点、难点）： Atomic & Molecular Orbitals Electrons surrounding atoms are concentrated into regions of space called atomic orbitals. The Heisenberg uncertainty principle states that it is impossible to know both the location and the momentum of an atomic particle, but it is possible to describe the probability that the electron will be found within a given region of space. The boundaries of an atomic orbital are commonly drawn to the region of 90% probability; there is a 90% probability that at any given time, the electron will be within the specified boundary. The electronic configuration of carbon is 1s2 2s2 2sp3. Atomic orbitals with s-character have spherical symmetry and a representation of the surface of the carbon 1s orbital is shown below. The wave properties of electrons make the description of the 2s orbital slightly more complex than the corresponding 1s orbital, in that, within the 2s sphere there is a region in which the amplitude of the electron standing wave falls to zero, that is, there is zero probability of finding the electron in this node region. Nodes are most easily seen in the description of the 2p atomic orbitals, which are shown below
2p atomic orbitals The electron densities along the x, y and z axes of the 2p orbitals are clearly shown in the figure; the nodes are the points at the origin and at these points, there is zero probability of finding the electron The sharing of electrons in a covalent bond occurs by overlap of the individual atomic orbitals. Head-on overlap between energetically compatible orbitals generates sigma (o)bonds, while sideways overlap (typically from adjacent pl orbitals)generates pi (T) bonds. Examples of sigma and T-bond bond formation between atoms a and"b are shown below B The nature of the bonding in hydrogen (H,)can be described using Molecular Orbital Theory. As the two ls atomic orbitals approach each other and begin to overlap, there is a decrease in the net energy of the system because the electrons in each atom tend to become attracted to the positive nucleus of the other atom, as well as its own nucleus. The more the orbitals overlap, the more the energy decreases, until the nuclei approach so closely that they begin to repel each. The point at which the repulsive and attractive forces balance defines the bond distance for a given covalent bond. 1s atomic 1s atomic bonding molecular orbital In molecular orbital theory, the number of atomic orbitals used to make the covalent bond must equal the total number of molecular orbitals in the molecule. In the example cited above, the atomic orbitals combine to form one bonding orbital, containing the two electrons, and one high-energy antibonding orbital which is empty. The molecular orbital description of this simple covalent
The electron densities along the x, y and z axes of the 2p orbitals are clearly shown in the figure; the nodes are the points at the origin and at these points, there is zero probability of finding the electron. The sharing of electrons in a covalent bond occurs by overlap of the individual atomic orbitals. Head-on overlap between energetically compatible orbitals generates sigma () bonds, while sideways overlap (typically from adjacent p orbitals) generates pi () bonds. Examples of sigma and -bond bond formation between atoms "A" and "B" are shown below. The nature of the bonding in hydrogen (H2) can be described using Molecular Orbital Theory. As the two 1s atomic orbitals approach each other and begin to overlap, there is a decrease in the net energy of the system because the electrons in each atom tend to become attracted to the positive nucleus of the other atom, as well as its own nucleus. The more the orbitals overlap, the more the energy decreases, until the nuclei approach so closely that they begin to repel each. The point at which the repulsive and attractive forces balance defines the bond distance for a given covalent bond. In molecular orbital theory, the number of atomic orbitals used to make the covalent bond must equal the total number of molecular orbitals in the molecule. In the example cited above, the atomic orbitals combine to form one bonding orbital, containing the two electrons, and one high-energy antibonding orbital which is empty. The molecular orbital description of this simple covalent
bonding is shown below. As described above, the bonding orbital is referred to as a orbital, while the corresponding antibonding orbital is referred to asad水- orbital Is ate u bonding molecular orbita In a similar manner, sideways overlap of adjacent p-orbitals forms a covalent Tt-orbital and a corresponding high-energy T-antibonding molecular orbital In general, electrons only populate antibonding orbitals when the molecule is in an excited state, and such orbitals are typically ignored in the discussion of organic reaction mechanisms. A more useful description of bonding is often given by the Valence Shell Electron Pair Repulsion(VSEPR) model, in which electrons are positioned around an atom to minimize electrostatic repulsion. This concept is described in more detail in the following section on orbital hybridization Bonding in Organic Molecules With an understanding of the tetrahedral nature of tetravalent carbon, organic compounds can be represented by a variety of structural drawings, as shown below
bonding is shown below. As described above, the bonding orbital is referred to as a -orbital, while the corresponding antibonding orbital is referred to as a *-orbital. In a similar manner, sideways overlap of adjacent p-orbitals forms a covalent -orbital and a corresponding high-energy *-antibonding molecular orbital. In general, electrons only populate antibonding orbitals when the molecule is in an excited state, and such orbitals are typically ignored in the discussion of organic reaction mechanisms. A more useful description of bonding is often given by the Valence Shell Electron Pair Repulsion (VSEPR) model, in which electrons are positioned around an atom to minimize electrostatic repulsion. This concept is described in more detail in the following section on orbital hybridization. Bonding in Organic Molecules With an understanding of the tetrahedral nature of tetravalent carbon, organic compounds can be represented by a variety of structural drawings, as shown below
The stick, or Deriding, model shows the carbon at the center of the tetrahedron (dark gray) with the hydrogens at each vertex (light gray); the covalent radius of each atom is approximated by the size of the color band. The ball-and-stick model provides similar information and is sometimes easier to visualize, and the true Van der Walls radius of the atoms is best shown by the space-filling model (shown with a ball-and-stick overlay). More complex hydrocarbons containing carbon chains can be formed by creating additional carbon-carbon bonds, as shown below for chains containing two, four and six carbon atoms you should note that in these structures, each carbon remains bonded to four other atoms (a valence of four) These molecules are also shown below in space-filling format s:+ When viewing organic molecules it is important to note that the rotation around carbon-carbon single bonds is generally very rapid (greater than 10. rotations per second)and the chain can assume a large number of conformations(termed conformational isomers)which are rapidly interconverted and cannot be separated under normal circumstances. A sample of conformational isomers for four- and six-carbon chains are shown below
The stick, or Deriding, model shows the carbon at the center of the tetrahedron (dark gray) with the hydrogens at each vertex (light gray); the covalent radius of each atom is approximated by the size of the color band. The ball-and-stick model provides similar information and is sometimes easier to visualize, and the true Van derWalls radius of the atoms is best shown by the space-filling model (shown with a ball-and-stick overlay). More complex hydrocarbons containing carbon chains can be formed by creating additional carbon-carbon bonds, as shown below for chains containing two, four and six carbon atoms; you should note that in these structures, each carbon remains bonded to four other atoms (a valence of four). These molecules are also shown below in space-filling format: When viewing organic molecules it is important to note that the rotation around carbon-carbon single bonds is generally very rapid (greater than 106 rotations per second) and the chain can assume a large number of conformations (termed conformational isomers) which are rapidly interconverted and cannot be separated under normal circumstances. A sample of conformational isomers for four- and six-carbon chains are shown below:
While drawings like those shown above most clearly show the structural features of organic molecules, it is clear that simpler methods are needed for routine representation of organic structures. The simplest, and least informative, is the simple representation of the molecular formula, showing ratios of atoms a alternative format in which all atoms and all covalent bonds are shown is the line-bond or Kekulee structure. While this does provide information regarding bonding in the molecule, structures of this type are tedious to draw, and organic molecules are more commonly represented using some form of abbreviated condensed structure CHaCH2CHCH3 CH2CH3 HH C6H14 H-C-C-C-C-C-H CH3 CH2CHCH2 CH3 CH3 CH2-CH-CH2-CH3 Line-Bond Molecular Formnla Assorted Condensed Formnlas Kekule A further condensation of structural information is accomplished in line" or structural drawings. In this format, each vertex in the drawing corresponds to a-ch group and each terminal line to a -Ch, group. It is assumed that all carbons have the appropriate number of hydrogens, and these are typically not shown. Multiple bonds are shown in structural drawings as double or triple lines, and it is again understood that the appropriate number of hydrogen atoms are attached CH3 Line or Structural Formulas
While drawings like those shown above most clearly show the structural features of organic molecules, it is clear that simpler methods are needed for routine representation of organic structures. The simplest, and least informative, is the simple representation of the molecular formula, showing ratios of atoms. A alternative format in which all atoms and all covalent bonds are shown is the "line-bond" or Kekuleé structure. While this does provide information regarding bonding in the molecule, structures of this type are tedious to draw, and organic molecules are more commonly represented using some form of abbreviated "condensed structure". A further condensation of structural information is accomplished in "line" or "structural" drawings. In this format, each vertex in the drawing corresponds to a -CH2- group and each terminal line to a -CH3 group. It is assumed that all carbons have the appropriate number of hydrogens, and these are typically not shown. Multiple bonds are shown in structural drawings as double or triple lines, and it is again understood that the appropriate number of hydrogen atoms are attached
In actual practice, you will find organic structures represented by hybrids of all of these methods (as shown above): line structures to represent the backbone of a molecule. and some variation of condensed structures to show side-chains, or three-dimensional information, etc. As a student of organic chemistry, it is essential to learn to recognize these structural representations and to be able to interchange formats and to use these to visualize the full three-dimensional molecule in question. Ionic, Covalent Polar bonding Bonding between atoms involves the transfer, or sharing, of electron density such that each atom is left with a stable outer shell of electrons. for atoms with significantly different electronegativity, the process of bonding generally involves complete electron transfer to form two species having net positive and negative charges. This type of bonding is termed ionic and simple salts are the classic examples (the electronegativities of sodium and chlorine are 0. 93 and 3. 16, respectively, and NaCl forms a fully ionized compound with distinct and separate sodium and chloride ions). For atoms with similar or identical electronegativities (such as chlorine, Cl, bonding involves the sharing of electron density between the two atoms, in a type of bonding which is termed covalent. While fully ionic and symmetrical covalent bonds represent the limits of bonding, the bonding in most organic (carbon-containing molecules falls somewhere between these two extremes. For example, carbon and chlorine have electronegativities of 2.55 and 3. 96, respectively. These values are close enough that the carbon-chlorine bond is covalent, but the bond is not symmetrical and the bulk of the electron density is associated with the hore electronegative chlorine. This polarization leaves a partial positive charge on the carbon and a matching partial negative charge on the chlorine. t is this polarization which gives organic compounds their chemical reactivity and allows simple structural units to be combined to form more complex
In actual practice, you will find organic structures represented by hybrids of all of these methods (as shown above); line structures to represent the backbone of a molecule, and some variation of condensed structures to show side-chains, or three-dimensional information, etc. As a student of organic chemistry, it is essential to learn to recognize these structural representations and to be able to interchange formats and to use these to visualize the full three-dimensional molecule in question. Ionic, Covalent & Polar Bonding Bonding between atoms involves the transfer, or sharing, of electron density such that each atom is left with a stable outer shell of electrons. For atoms with significantly different electronegativity, the process of "bonding" generally involves complete electron transfer to form two species having net positive and negative charges. This type of bonding is termed ionic and simple salts are the classic examples (the electronegativities of sodium and chlorine are 0.93 and 3.16, respectively, and NaCl forms a fully ionized compound with distinct and separate sodium and chloride ions). For atoms with similar or identical electronegativities (such as chlorine, Cl2) bonding involves the sharing of electron density between the two atoms, in a type of bonding which is termed covalent. While fully ionic and symmetrical covalent bonds represent the limits of bonding, the bonding in most organic (carbon-containing) molecules falls somewhere between these two extremes. For example, carbon and chlorine have electronegativities of 2.55 and 3.96, respectively. These values are close enough that the carbon-chlorine bond is covalent, but the bond is not symmetrical and the bulk of the electron density is associated with the more electronegative chlorine. This polarization leaves a partial positive charge on the carbon and a matching partial negative charge on the chlorine. It is this polarization which gives organic compounds their chemical reactivity and allows simple structural units to be combined to form more complex
molecules, i.e., the science of organic synthesis partial positive partial negative Increased electron density associated with the electronegative ch Ionic structures are often represented using Lewis dot formulas, in which the electrons in the outer shell are shown as paired or unpaired dots surrounding the atomic symbol. Covalent bonds can likewise be shown using dot formulas and these are often useful. The large number of bonds in a typical organic molecule, however, would make drawing lewis dot structures somewhat tedious and covalent bonds are typically represented as a single line connecting the two atoms. When it is desirable, polarization along a covalent bond can be shown in these structures by using the symbols 8+ and8- over the atoms in question, as shown below. se Cl Cl Cl 8 Cl: Cl: C: Cl: C1-C-Cl &eCI-C-C18e Cl: se 教学手段与方法:课堂讲授多媒体教学 思考题、讨论题、作业: Homeworks;Page312,313,3.14,3.15,3.17,3.18.3.23, 3.24. 参考资料(含参考书、文献等) 1. Solomons, Organic Chemistry, fifth adition 2. Oxford; Organic Chemistr 3.北京大学,有机化学 4.南京大学,有机化学,(上,下) 5.邢其毅,有机化学,(上,下) 注:1、每项页面大小可自行添减;2一次课为一个教案;3、“重点"、“难点"、“教学手段
molecules, i.e., the science of organic synthesis. Ionic structures are often represented using Lewis dot formulas, in which the electrons in the outer shell are shown as paired or unpaired "dots" surrounding the atomic symbol. Covalent bonds can likewise be shown using dot formulas, and these are often useful. The large number of bonds in a typical organic molecule, however, would make drawing Lewis dot structures somewhat tedious, and covalent bonds are typically represented as a single line connecting the two atoms. When it is desirable, polarization along a covalent bond can be shown in these structures by using the symbols "+" and "-" over the atoms in question, as shown below. 教学手段与方法：课堂讲授.多媒体教学。 思考题、讨论题、作业：Homeworks;Page.3.12, 3.13, 3.14, 3.15, 3.17, 3.18.3.23, 3.24. 参考资料（含参考书、文献等）： 1. Solomons, Organic Chemistry, fifth adition 2. Oxford; Organic Chemistry 3. 北京大学， 有机化学 4．南京大学， 有机化学，（上，下） 5．邢其毅，有机化学， （上，下） 注：1、每项页面大小可自行添减；2 一次课为一个教案；3、“重点”、“难点”、“教学手段
与方法”部分要尽量具体;4、授课类型指:理论课、讨论课、实验或实习课、练习或习题 课等
与方法”部分要尽量具体；4、授课类型指：理论课、讨论课、实验或实习课、练习或习题 课等
