专业英语.docx

1、专业英语Chapter I Structure & Bonding The study of organic chemistry must at some point extend to the molecular level, for the physical and chemical properties of a substance are ultimately explained in terms of the structure and bonding of molecules. This module 模块introduces some basic facts and princi

2、ples that are needed for a discussion of organic molecules. 1-1 Electronic ConfigurationsElectron Configurations in the Periodic Table The periodic table shown here is severely truncated. There are, of course, over eighty other elements. The halogens are one electron short of a valence shell octet,

3、and are among the most reactive of the elements. In their chemical reactions halogen atoms achieve a valence shell octet by capturing or borrowing the eighth electron from another atom or molecule. The alkali metals are also exceptionally reactive, but for the opposite reason. These atoms have only

4、one electron in the valence shell, and on losing this electron arrive at the lower shell valence octet. As a consequence of this electron loss, these elements are commonly encountered as cations (positively charged atoms).1-2 Chemical Bonding and Valence 1-2-1 Ionic Bonding When sodium is burned in

5、a chlorine atmosphere, it produces sodium chloride. This has a high melting point (800 C) and dissolves in water to give a conducting solution. Sodium chloride is an ionic compound, in which an electron of sodium atom was transferred to a chlorine atom and generates a sodium cation and a chloride an

6、ion. Electrostatic attraction results in these oppositely charged ions packing together in a lattice. The attractive forces holding the ions in place can be referred to as ionic bonds. 1-2-2 Covalent Bonding Water is a liquid at room temperature; carbon dioxide and carbon tetrafluoride are gases. No

7、ne of these compounds is composed of ions. A different attractive interaction between atoms, called covalent bonding, is involved here. Covalent bonding occurs by sharing of valence electrons, rather than an outright electron transfer. Similarities in physical properties (they are all gases) suggest

8、 that the diatomic elements H2, N2, O2, F2 & Cl2 also have covalent bonds. Carbon dioxide is notable because it is a case in which two pairs of electrons are shared by the same two atoms. This is an example of a double covalent bond.1-2-3 Valence The number of electrons an atom gain or lose to achie

9、ve a valence octet is called valence. The valences here represent the most common form 普通形式 in organic compounds. Many elements, such as chlorine, bromine and iodine, are known to exist in several valence states in different inorganic compounds. 1-3 Charge Distribution If the electron pairs in coval

10、ent bonds were shared absolutely evenly there would be no fixed local charges within a molecule. Although this is true for diatomic elements such as H2, N2 and O2, most covalent compounds show some degree of local charge separation, resulting in bond and / or molecular dipoles. 1-3-1 electronegativi

11、ty Different atoms have different affinities for nearby electrons. The ability of an element to attract or hold onto electrons is called electronegativity. Fluorine has the greatest electronegativity of all the elements, and the heavier alkali metals such as potassium, rubidium and cesium have the l

12、owest electronegativities. It should be noted that carbon is about in the middle of the electronegativity range, and is slightly more electronegative than hydrogen. 1-3-2 Polar Covalent Bonds When two different atoms are bonded covalently, the shared electrons are attracted to the more electronegati

13、ve atom of the bond, resulting in a shift of electron density toward the more electronegative atom. Such a covalent bond is polar, and will have a dipole. The degree of polarity and the magnitude of the bond dipole will be proportional to the difference in electronegativity of the bonded atoms. Thus

14、 a OH bond is more polar than a CH bond, with the hydrogen atom of the former being more positive than the hydrogen bonded to carbon. Likewise, CCl and CLi bonds are both polar, but the carbon end is positive in the former and negative in the latter. The dipolar nature of these bonds is often indica

15、ted by a partial charge notation (+/) or by an arrow pointing to the negative end of the bond. The shift of electron density in a covalent bond toward the more electronegative atom or group can be observed in several ways. For bonds to hydrogen, acidity is one criterion. If the bonding electron pair

16、 moves away from the hydrogen nucleus the proton will be more easily transfered to a base (it will be more acidic). Methane is almost non-acidic, since the CH bond is nearly non-polar. The OH bond of water is polar, and it is at least 25 powers of ten more acidic than methane. HF is over 12 powers o

17、f ten more acidic than water as a consequence of the greater electronegativity difference in its atoms. Electronegativity differences may be transmitted through connecting covalent bonds by an inductive effect. This inductive transfer of polarity tapers off as the number of transmitting bonds increa

18、ses, and the presence of more than one highly electronegative atom has a cumulative effect. For example, trifluoro ethanol, CF3CH2OH is about ten thousand times more acidic than ethanol, CH3CH2OH. 1-3-2 Functional Groups Functional groups are atoms or small groups of atoms (two to four) that exhibit

19、 a characteristic reactivity when treated with certain reagents. A particular functional group will almost always display its characteristic chemical behavior when it is present in a compound. Because of their importance in understanding organic chemistry, functional groups have characteristic names

20、 that often carry over in the naming of individual compounds incorporating specific groups. 1-4 The Shape of Molecules The three dimensional shape or configuration of a molecule is an important characteristic. Three dimensional configurations are best viewed with the aid of models. In order to repre

21、sent such configurations on a two-dimensional surface (paper, blackboard or screen), we often use perspective drawings in which the direction of a bond is specified by the line connecting the bonded atoms. A simple straight line represents a bond lying approximately in the surface plane. The two bon

22、ds to substituents A in the structure on the left are of this kind. A wedge shaped bond is directed in front of this plane (thick end toward the viewer), as shown by the bond to substituent B; and a hatched bond is directed in back of the plane (away from the viewer), as shown by the bond to substit

23、uent D. The following examples make use of this notation, and also illustrate the importance of including non-bonding valence shell electron pairs (colored blue) when viewing such configurations . Bonding configurations are readily predicted by valence-shell electron-pair repulsion theory, commonly

24、referred to as VSEPR in most introductory chemistry texts. This simple model is based on the fact that electrons repel each other, and that it is reasonable to expect that the bonds and non-bonding valence electron pairs will prefer to be as far apart as possible. The bonding configurations of carbo

25、n are easy to remember, since there are only three categories. In the three examples shown above, the central atom (carbon) does not have any non-bonding valence electrons; consequently the configuration may be estimated from the number of bonding partners alone. For molecules of water and ammonia,

26、however, the non-bonding electrons must be included in the calculation. In each case there are four regions of electron density associated with the valence shell so that a tetrahedral bond angle is expected. The measured bond angles of these compounds (H2O 104.5 & NH3 107.3) show that they are close

27、r to being tetrahedral than trigonal or linear. Of course, it is the configuration of atoms (not electrons) that defines the the shape of a molecule, and in this sense ammonia is said to be pyramidal (not tetrahedral). The compound boron trifluoride, BF3, does not have non-bonding valence electrons

28、and the configuration of its atoms is trigonal. 1-5 Isomers1-5-1 Structural Formulas It is necessary to draw structural formulas for organic compounds because in most cases a molecular formula does not uniquely represent a single compound. Different compounds having the same molecular formula are ca

29、lled isomers . When the group of atoms that make up the molecules of different isomers are bonded together in fundamentally different ways, we refer to such compounds as constitutional isomers. There are seven constitutional isomers of C4H10O, and structural formulas for these are drawn in the follo

30、wing table. These formulas represent all known and possible C4H10O compounds, and display a common structural feature. There are no double or triple bonds and no rings in any of these structures. Simplification of structural formulas may be achieved without any loss of the information they convey. I

31、n condensed structural formulas the bonds to each carbon are omitted, but each distinct structural unit (group) is written with subscript numbers designating multiple substituents, including the hydrogens. Shorthand (line) formulas omit the symbols for carbon and hydrogen entirely. Each straight lin

32、e segment represents a bond, the ends and intersections of the lines are carbon atoms, and the correct number of hydrogens is calculated from the tetravalency of carbon. Non-bonding valence shell electrons are omitted in these formulas. 1-5-2 Distinguishing Carbon Atoms When discussing structural formulas, it is often useful to distinguish different groups of carbon atoms by their structural characteristics. A primary carbon (1) is one that is bonded to no more than one other carbon atom. A secondary carbon (2) is bonded to two