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: