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Chemical Reactivity Organic chemistry encompasses a very large number of compounds ( many millions ), and our previous discussion and illustrations have focused on their structural characteristics. Now that we can recognize these actors ( compounds ), we turn to the roles they are inclined to play in the scientific drama staged by the multitude of chemical reactions that define organic chemistry.  We begin by defining some basic terms that will be used frequently as this subject is elaborated.                                                   Chemical Reaction A transformation resulting in a change of composition, constitution and/or configuration of a compound ( referred to as the reactant or substrate ).                                               Reactant or Substrate  The organic compound undergoing change in a chemical reaction. Other compounds may also be involved, and common reactive partners ( reagents ) may be identified. The reactant is often ( but not a

Water has been referred to as the "universal solvent", and its widespread distribution on this planet and essential role in life make it the benchmark for discussions of solubility. Water dissolves many ionic salts thanks to its high dielectric constant and ability to solvate ions. The former reduces the attraction between oppositely charged ions and the latter stabilizes the ions by binding to them and delocalizing charge density. Many organic compounds, especially alkanes and other hydrocarbons, are nearly insoluble in water. Organic compounds that are water soluble, such as most of those listed in the above table, generally have hydrogen bond acceptor and donor groups. The least soluble of the listed compounds is diethyl ether, which can serve only as a hydrogen bond acceptor and is 75% hydrocarbon in nature. Even so, diethyl ether is about two hundred times more soluble in water than is pentane. The chief characteristic of water that influences these solubilities is

Most organic compounds have melting points below 200 ºC. Some decompose before melting, a few sublime, but a majority undergo repeated melting and crystallization without any change in molecular structure. When a pure crystalline compound is heated, or a liquid cooled, the change in sample temperature with time is roughly uniform. However, if the solid melts, or the liquid freezes, a discontinuity occurs and the temperature of the sample remains constant until the phase change is complete. This behavior is shown in the diagram on the right, with the green segment representing the solid phase, light blue the liquid, and red the temperature invariant liquid/solid equilibrium. For a given compound, this temperature represents its melting point (or freezing point), and is a reproducible constant as long as the external pressure does not change. The length of the horizontal portion depends on the size of the sample, since a quantity of heat proportional to the heat of fusion must be

The most powerful intermolecular force influencing neutral (uncharged) molecules is the   hydrogen bond . If we compare the boiling points of methane (CH 4 ) -161ºC, ammonia (NH 3 ) -33ºC, water (H 2 O) 100ºC and hydrogen fluoride (HF) 19ºC, we see a greater variation for these similar sized molecules than expected from the data presented above for polar compounds. This is shown graphically in the following chart. Most of the simple hydrides of group IV, V, VI & VII elements display the expected rise in boiling point with molecular mass, but the hydrides of the most electronegative elements (nitrogen, oxygen and fluorine) have abnormally high boiling points for their mass. The exceptionally strong dipole-dipole attractions that cause this behavior are called the  hydrogen bond . Hydrogen forms polar covalent bonds to more electronegative atoms such as oxygen, and because a hydrogen atom is quite small, the positive end of the bond dipole (the hydrogen) can approach neighbor

The molecule is the smallest observable group of uniquely bonded atoms that represent the composition, configuration and characteristics of a pure compound. Our chief focus up to this point has been to discover and describe the ways in which atoms bond together to form molecules. Since all observable samples of compounds and mixtures contain a very large number of molecules ( ca. !0 20 ), we must also concern ourselves with interactions between molecules, as well as with their individual structures. Indeed, many of the physical characteristics of compounds that are used to identify them (e.g. boiling points, melting points and solubilities) are due to intermolecular interactions. All atoms and molecules have a weak attraction for one another, known as  van der Waals  attraction. This attractive force has its origin in the electrostatic attraction of the electrons of one molecule or atom for the nuclei of another. If there were no van der Waals forces, all matter would exist in a

A more detailed model of covalent bonding requires a consideration of valence shell atomic orbitals. For second period elements such as carbon, nitrogen and oxygen, these orbitals have been designated 2 s , 2p x , 2p y   & 2p z . The spatial distribution of electrons occupying each of these orbitals is shown in the diagram below.   Very nice displays of orbitals may be found at the following sites: J. Gutow, Univ. Wisconsin Oshkosh R. Spinney, Ohio State  M. Winter, Sheffield University The valence shell electron configuration of carbon is 2 s 2 , 2p x 1 , 2p y 1  & 2p z 0 . If this were the configuration used in covalent bonding, carbon would only be able to form two bonds. In this case, the valence shell would have six electrons- two shy of an octet. However, the tetrahedral structures of methane and carbon tetrachloride demonstrate that carbon can form four equivalent bonds, leading to the desired octet. In order to explain this covalent bonding, Linus Pauling prop

Kekulé structural formulas are essential tools for understanding organic chemistry. However, the structures of some compounds and ions cannot be represented by a single formula. For example, sulfur dioxide (SO 2 ) and nitric acid (HNO 3 ) may each be described by two equivalent formulas (equations 1 & 2). For clarity the two ambiguous bonds to oxygen are given different colors in these formulas. 1)   sulfur dioxide 2)   nitric acid If only one formula for sulfur dioxide was correct and accurate, then the double bond to oxygen would be shorter and stronger than the single bond. Since experimental evidence indicates that this molecule is bent (bond angle 120º) and has equal length sulfur : oxygen bonds (1.432 Å), a single formula is inadequate, and the actual structure resembles an average of the two formulas. This averaging of electron distribution over two or more hypothetical contributing structures ( canonical forms ) to produce a hybrid electronic structure is

Although structural formulas are essential to the unique description of organic compounds, it is interesting and instructive to evaluate the information that may be obtained from a  molecular formula alone. Three useful rules may be listed: The number of hydrogen atoms that can be bonded to a given number of carbon atoms is limited by the valence of carbon . For compounds of carbon and hydrogen (hydrocarbons) the maximum number of hydrogen atoms that can be bonded to n carbons is  2n + 2  (n is an integer). In the case of methane, CH 4 ,  n=1  &  2n + 2 = 4 . The origin of this formula is evident by considering a hydrocarbon made up of a chain of carbon atoms. Here the middle carbons will each have two hydrogens and the two end carbons have three hydrogens each. Thus, a six-carbon chain (n = 6) may be written H-(CH 2 ) 6 -H, and the total hydrogen count is  (2 x 6) + 2 = 14 . The presence of oxygen (valence = 2) does not change this relationship, so the previously described

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 called  isomers , and the prevalence of organic isomers reflects the extraordinary versatility of carbon in forming strong bonds to itself and to other elements. 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 C 4 H 10 O, and structural formulas for these are drawn in the following table. These formulas represent all known and possible C 4 H 10 O compounds, and display a common structural feature.  There are no double or triple bonds and no rings in any of these structures . . Note that each of the carbon atoms is bonded to four other atoms, and is  saturated  with bonding part

 The three dimensional shape or   configuration   of a molecule is an important characteristic. This shape is dependent on the preferred spatial orientation of covalent bonds to atoms having two or more bonding partners. Three dimensional configurations are best viewed with the aid of models.   In order to represent 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.   In most cases the focus of configuration is a carbon atom so the lines specifying bond directions will originate there. As defined in the diagram on the right, a simple straight line represents a bond lying approximately in the surface plane. The two bonds 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