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F. Richard Daley and Sally J. Daley www.ochem4free.com Organic Chemistry Chapter 4 Physical Properties of Organic Compounds 4.1 Phases of Matter 175 Sidebar - Liquid Crystals 4.2 Melting Points 179 4.3 Boiling Points 183 4.4 Solubility 190 Sidebar - Surfactants 4.5 Density 197 Key Ideas from Chapter 4 178 194 199 Organic Chemistry - Ch 4 173 Daley & Daley Copyright 1996-2005 by Richard F. Daley & Sally J. Daley All Rights Reserved. No part of this publication may be reproduced, stored in a retrieval system, or transmitted in any form or by any means, electronic, mechanical, photocopying, recording, or otherwise, without the prior written permission of the copyright holder. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 174 Daley & Daley Chapter 4 Physical Properties of Organic Compounds Chapter Outline 4.1 4.2 4.3 4.4 4.5 Phases of Matter A review of the common phases of matter: solid, liquid, and gas Melting and Freezing Points Factors affecting the melting and freezing points of a compound Boiling Points Factors affecting the boiling point of a compound Solubility Why a compound is soluble in a given solvent Density Factors affecting the density of a compound Objectives Understand what the terms solid, liquid, and gas mean at the molecular level Know how the symmetry and conformation of an organic structure affect its melting point Understand how a functional group affects the melting point of a compound Know how structure and functional groups affect the boiling point of a compound Recognize the molecular features that affect the solubility of a compound Know what molecular features affect the density of a compound Be able to rank compounds in order based on the physical properties discussed www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 175 Daley & Daley Contrariwise, continued Tweedledee, if it was so, it might be; and if it were so, it would be: but as it isn't, it ain't. That's logic. Lewis Carroll he physical properties of a compound include such things as its color, odor, refractive index, density, solubility, melting point, and boiling point. The exact values of the physical properties of a compound depend on its molecular structure. All the physical properties of a compound correlate with its molecular structure. Although chemists cannot examine the structure of a compound and predict the exact values for these physical properties, they can look at two compounds with similar structures and determine which has the higher melting or boiling point. By knowing the structure of a compound, a chemist can also choose a better solvent to extract or recrystallize the compound. Knowledge of physical propertiessuch as solubilities, melting or boiling points, and densitywill assist you in making judgments about the practical aspects of the isolation and purification of new molecules. This translates into doing better laboratory work. Understanding how molecules interact physically leads to a better understanding of how they interact in a chemical reaction. This chapter shows you how to use the molecular structure to predict qualitatively the physical properties of compounds. You can qualitatively predict physical properties because they relate to differences in intermolecular forces. T 4.1 Phases of Matter Matter ordinarily exists in four phases: solid, liquid, gas, and plasma. Plasma is a high-temperature phase not encountered in the typical organic chemistry laboratory and thus not considered in this book. Most pure chemical compounds exist in each of these phases. The phase of a compound at a particular moment depends on its temperature and pressure. For example, water at 1 atmosphere pressure is a solid below 0oC, a liquid from 0oC to 100oC, and a gas above 100oC. The temperatures at which a compound makes the transitions between phases are unique for each compound. These temperatures are a measure of some of the unique physical properties of that compound. Because you have studied solids, liquids, and gases from your earliest years of school, you already know many facts about them. The goal for this section is to look at solids, liquids, and gases from the www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 176 Daley & Daley Intermolecular interactions are interactions between molecules. In the crystalline form the molecules pack together into a threedimensional matrix. In the amorphous form the molecules pack together in an irregular pattern. An ordered matrix is a regular array of particles with each particle surrounded by the same arrangement of particles as any other particle. organic chemistry point of viewto cover their characteristics and to define any terms that are important in organic chemistry. Consider a hypothetical ideal compound as it makes its transition through each phase. For simplicity's sake, assume that each molecule of this compound is spherical. A sphere is a simple shape that is easy to visualize with minimal intermolecular interactions (interactions between molecules) Starting at some low temperature, the compound is a solid. There are two extremes of solid forms: the crystalline form and the amorphous form. Because nearly all solid organic compounds are in the crystalline form, this book does not discuss the amorphous form. Spheres pack easily into an ordered matrix, thus, the idealized solid is a crystal. Envision this crystalline form by considering a box full of identically sized marbles. The boxs size is such that one layer of marbles fits exactly in the bottom. The consecutive layers of marbles then fit on top of each other in an ordered set of layers. Figure 4.1 illustrates such a regular ordered matrix of molecules in a crystalline solid. Figure 4.1. Idealized spherical molecules ordered in a very regular fashion in the crystalline form of the solid phase. Fusion is the transition between the solid and liquid phases. The melting point of a compound is the temperature at which fusion occurs. The intermolecular attractive forces hold the molecules in place in the crystal lattice. The stronger these forces are, the greater the energy that is required to break down the crystal lattice. When you apply sufficient heat to a solid, the crystalline structure of that solid breaks down. As you heat a solid, it undergoes a transition from the solid phase to the liquid phase. Chemists call this transition fusion and the characteristic temperature at which a compound undergoes fusion its melting point. In the liquid phase the molecules are no longer packed in an ordered matrix. There is more disorder, but still a definite boundary between the liquid and its surroundings. Figure 4.2 illustrates the liquid phase using the idealized compound. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 177 Daley & Daley Figure 4.2. The liquid phase has less order than the solid phase. The organization and orientation of the molecules are more random. Vaporization is the transition between the liquid and gas phases. The boiling point of a compound is the temperature at which vaporization occurs. As you continue to increase the temperature, the compound undergoes another transformation. It moves from its liquid phase to its gas phase. Chemists call the process of going from the liquid phase to the gas phase vaporization. The characteristic temperature at which a compound undergoes vaporization is its boiling point. In its gas phase a compound has even more disorder than it has in its liquid phase. There is little intermolecular interaction because the molecules now have enough energy to counteract these forces. Since there is so little intermolecular interaction, the molecules move freely, filling all the available space. There is no obvious boundary between the molecules of the gas and its surroundings. Figure 4.3 illustrates the gas phase using the idealized compound. Figure 4.3. The gas phase is in even more disorder than the liquid phase. The molecules have little intermolecular interaction. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 178 Daley & Daley [Sidebar ] Liquid Crystals A liquid crystal is a material showing characteristics of both a liquid and a crystal. A mesophase shows behavior that is intermediate between two phases. In a liquid crystal, the individual molecules have restricted mobility. They are more ordered than in a true liquid but less ordered than in a crystalline solid; thus, they are in an intermediate state. Because of this intermediate state, the liquid crystalline state is often called a mesophase. In general, materials that exhibit liquid crystallinity have rigid rod- or disk-like shapes. O CH3CH2CH2OCO O CO OCH2CH3 Liquid crystal between 55oC and 87oC With nematic ordering, the molecules order themselves along their long axes. With smectic ordering, not only do the molecules order themselves along their long axes, but they also have a layering tendency. A liquid crystal contains many small regions, each with a million or so molecules, with a preferred orientation for the molecules. Thus, even though the molecules in a liquid are in constant motion, the molecules in a particular region are all aligned with each other. On a larger scale, many regions of preferred orientation exist, but these regions are randomly oriented with respect to one another. Liquid crystals are found in two types of ordering: the nematic form and the smectic form. Figure 4.4 illustrates both types of ordering. Weak intermolecular forces produce nematic ordering, but stronger intermolecular forces produce smectic ordering. Stronger intermolecular forces also produce larger regions of order in the liquid. (a) Nematic form of a liquid crystalline material. (b) Smectic form of a liquid crystalline material. Figure 4.4. Schematic comparison between (a) the nematic form and (b) the smectic form of a liquid crystalline material. An example of a liquid crystal is 4-octyl-4-cyanobiphenyl. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 179 Daley & Daley CH3CH2CH2CH2CH2CH2CH2CH2 CN 4'-Octyl-4-cyanobiphenyl Liquid crystal at room temperature At room temperature it adopts the smectic form. Figure 4.5 is a scanning tunneling microscope image clearly showing the smectic alignment of the molecules. (NEED PERMISSION FOR FIGURE 1A FROM REF BELOW) Figure 4.5. A scanning tunneling microscope image of the ordering of the molecules of 4-octyl-4-cyanobiphenyl. (from Smith and Horber, et.al., Science, 7/7/89, p43) Liquid crystals are widely used in electronic displays because an external force, such as an electric field, increases the size of the ordered regions. Liquid crystals with large ordered regions have different optical properties than liquid crystals with small ordered regions. Most significantly, the large ordered regions become visible. Some watches and computer displays have a sealed container of liquid crystals. The surface of the container has a pattern of transparent electrodes. Depending on which electrodes are powered, a pattern of dots, lines, or symbols appears comprising the information visible on the display. 4.2 Melting Points Melting occurs as a result of applying sufficient heat to a solid so that it moves from its solid phase to its liquid phase. The temperature at which this transformation takes place is the compounds melting point. Each compound has a specific melting point; thus, when working with an unknown solid compound or a known solid compound whose purity is in doubt, one important measurement that chemists often take is the compound's melting point. Fortunately, taking a melting point is relatively easy to do. They then use this temperature to help identify the unknown solid compound or to verify that the compound is impure. The melting point of a compound is the temperature at which the compounds solid and liquid phases are in equilibrium when the temperature is being increased. Freezing Points www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 180 Daley & Daley The freezing point of a compound is the temperature at which the compounds solid and liquid phases are in equilibrium when the temperature is being decreased. Not only does a compound have a melting point, it also has a freezing point. The freezing point of a compound is the temperature at which a compounds solid and liquid phases are in equilibrium when the compound is moving from its liquid phase to its solid phase. (Note that the direction of temperature change for freezing is opposite to the direction of temperature change for melting.) The melting point and the freezing point for a compound are the same temperature. Although freezing is a process that commonly occurs, chemists seldom use freezing points because controlling the cooling process is more difficult than controlling the heating process. Chemists do use freezing points for compounds that freeze at or near room temperature. When working with a liquid compound, they usually find its boiling point. Section 4.3 discusses boiling points. Four factors that influence a compounds melting point are symmetry, polarity, hydrogen bonding, and molecular weight. Before looking at those factors, here is some information that you need to keep in mind about solids. As you learned in Section 4.1, a crystalline solid is composed of molecules arranged in a regular pattern. To arrange themselves in this pattern, the molecules normally have the same conformation and the same specific orientation relative to each other. Figure 4.6 shows a hypothetical crystalline solid. Above the solid are molecules A and B. They are both the same as the molecules in the solid, but neither molecule A nor molecule B can fit into the crystalline lattice. Molecule A has an incorrect conformation; and although molecule B has the correct conformation, it has an incorrect orientation. A B Figure 4.6. Molecule A has the wrong conformation to fit into the crystalline structure. Molecule B has the wrong orientation. A symmetrical molecule has a regular shapespherical, cubical, cylindrical, etc. The more symmetrical a molecule is, the better it fits into a crystalline structure, and the higher is its melting point. Because molecules must have a similar conformation to form a crystalline lattice, any molecules with the wrong symmetry do not easily fit into the growing crystal. For example, consider the following isomeric www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 181 Daley & Daley compoundspentane with a melting point of 130oC and 2,2dimethylpropane with a melting point of 17oC. Pentane exists in a large variety of conformations that are similar in energy to each other. Below are three of its many possible staggered conformations. H H3C H H H H CH3 H H3C H H H H H H CH3 H3C H H H H H H CH3 Some conformations of pentane Because the rotational energy barrier for the various conformations of pentane is very low, pentane exists as a mixture of conformations except at a very low temperature. Thus, pentane does not readily form a solid. It favors the liquid phase over the solid phase in the solidliquid equilibrium. In contrast to the conformational mobility of pentane, 2,2dimethylpropane has only one conformation of the carbon skeleton: H3C H3C CH3 CH3 C 2,2-Dimethylpropane m.p. -17oC In a polar molecule one or more bonds exhibit bond dipoles, and generally the molecule has a dipole moment. All the hydrogen atoms of 2,2-dimethylpropane are equivalent, and the molecule has a very high level of symmetry. This symmetry means that each molecule has the right conformation to form a crystal. Also, this symmetry increases the probability that each molecule has the correct orientation to fit into a growing crystal. Thus, 2,2dimethylpropane readily forms a crystalline lattice. It favors its solid form in the solid-liquid equilibrium and requires a higher temperature than pentane to convert from its solid form to its liquid form. Another important factor contributing to the melting point of a compound is its polarity. The more polar a compound is, the stronger are its intermolecular attractions and the higher is its melting point. For example, benzyl alcohol melts at 15oC whereas benzoic acid melts at 122oC. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 182 Daley & Daley O CH2OH COH Benzyl alcohol m.p. -15oC Benzoic acid m.p. 122oC A dimer is an association between two molecules that are held together by hydrogen bonds. Hydrogen bonds form between polar oxygens in one molecule and acidic hydrogens in another. Section 4.3, page 000 discusses them in more detail. Both compounds are polar, but the intermolecular attractions of benzoic acid are much stronger than those of benzyl alcohol, as is shown by the fact that benzoic acid forms a relatively stable dimer. This dimer is held together by attractions, called hydrogen bonds, between the polar oxygens and the acidic hydrogens of the carboxylic acid functional groups. Hydrogen bonds influence the melting point of a solid, but they influence the boiling point of a liquid even more (see Section 4.3.) O C O H O H O C Dimer of benzoic acid The molecular weight of a molecule is the sum of the atomic weights of all of the atoms in that molecule. The molecular weight of a compound influences its melting point. As the molecular weight increases in a homologous series of compounds, so does the melting point. The reason the melting point increases with the weight is that it takes more energy to separate larger molecules from a crystalline structure than it takes to separate smaller ones. Although the increase in melting point holds true over a long series, other factors at times override the effects of molecular weight for molecules of similar size. Two of these factors are symmetry and hydrogen bonding. For example, because methane molecules are symmetrical and pack into a crystalline lattice better than propane molecules do, methanes melting point (182.5 oC) is seven degrees higher than propane's (189.7 oC) despite its lower molecular weight. Exercise 4.1 a) One stereoisomer of 1,4-dibromocyclohexane has a melting point of 113oC; the other isomer melts below 0oC. Draw both structures and assign melting points to the structures. Justify your choice. b) Which has the higher melting point, hexane or cyclohexane? Explain your choice. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 183 Daley & Daley c) Cyclopentanethiol has a markedly lower melting point than cyclopentanol even though it has a higher molecular weight. Explain. SH Cyclopentanethiol OH Cyclopentanol Sample solution b) Cyclohexane has the higher melting point because it is more symmetrical and packs more readily into a crystal lattice. Hexane has many more conformations and thus is less likely to fit into a crystalline structure. Cyclohexane Hexane 4.3 Boiling Points An important physical property of a liquid is its boiling point, which is the transformation point of a compound from its liquid phase to its gas phase. The definition of boiling point is the temperature at which the vapor pressure of a liquid equals the external (atmospheric) pressure above that liquid. When considering the boiling points for a homologous series of molecules, molecular weight is an important factor. Table 4.1 lists the boiling points for several common homologous series. As you look at the table, notice that within a series the boiling points quickly increase. For each series, the boiling point shows a fairly regular increase of 20-30oC with each additional CH2 group in the chain. Also, the boiling points vary drastically based on the functional group that the molecule contains. Boiling Points (oC) RCl ROH 24 12 47 65 78 97 The vapor pressure of a compound is the pressure exerted by the vapor above the compounds surface. Number of Carbons in R 1 2 3 RH 160 89 42 RCOOH 101 118 141 www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 184 Daley & Daley Number of Carbons in R 4 5 6 7 8 9 RH 0 36 69 98 126 151 Boiling Points (oC) RCl ROH 78 98 134 160 183 203 117 138 157 176 195 213 RCOOH 164 186 206 223 240 254 Table 4.1. Comparative boiling points for several selected homologous series. A complementary polarization occurs when a temporary dipole in one molecule induces a similar dipole in another molecule. The van der Waals forces, also called London dispersion forces, are intermolecular attractions arising from complementary polarizations. The molecules in the liquid phase are still close enough to each other that their physical interactions are similar to the physical interactions that occur between the same molecules in their solid phase. Although the interactions in the liquid phase are far more random, and thus generally weaker, than in the solid phase, they still occur and are important. In the gas phase, the molecules are so far apart that the intermolecular interactions are much weaker and are generally of little importance. Except for highly polar substances, such as carboxylic acids, there are essentially no interactions between most molecules in the gas phase. The interactions between molecules are the result of attractions between those molecules. These attractions fit into three different categories: van der Waals forces, dipolar attractions, and hydrogen bonding. The energies associated with these interactions are small compared to those associated with chemical bonds, but for a collection of molecules, they are significant. When a pair of molecules approach each other, the nonbonded electrons on one molecule tend to attract the partially positive atoms on the other molecule. These attractive forces increase until they reach a maximum at intermolecular distances between 200 and 400 pm. At distances closer than this, the molecules tend to repel one another because their electrons repel one another. The actual distances at which the molecules begin to repel one another is the sum of the van der Waals radii of the two groups. The average distance between molecules in the liquid phase is in the range of 200-400 pm. To understand what happens with these attractive forces, consider two nonpolar molecules, X and Y. Keep in mind that the distribution of electrons in these molecules is continually fluctuating. As the two molecules approach each other, they experience a mutual attraction, and any polarization in molecule X induces a complementary polarization in molecule Y. Chemists call this attraction van der Waals forces or London dispersion forces. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 185 Daley & Daley + -+ X van der Waals interactions ++ Y -+ +X -+ +Y van der Waals forces, which are sometimes called induced polarizations or induced dipoles, are only temporary and are constantly changing because the electron distribution within each molecule rapidly fluctuates. When the polarization in one molecule changes, it influences a neighboring molecule, which in turn influences another neighboring molecule. The net effect is that all neighboring molecules are attracted to each other. The magnitude of van der Waals forces is based on the number of electrons in the molecules and how many of those electrons participate in these induced dipole-dipole interactions. For a low polarity liquid to boil, it must overcome the van der Waals forces. The major factor in the magnitude of these forces is the shape of the molecule. Highly branched molecules have a more spherical shape and smaller van der Waals attractions. Unbranched molecules have more surface area that can be involved in intermolecular interactions and higher van der Waals attractions because they can pack closer. You can see this effect in the boiling points of the following three isomers: pentane, 2-methylbutane, and 2,2-dimethylpropane. CH3 CH3CH2CH2CH2CH3 CH3CH2CHCH3 CH3 Pentane b.p. 36oC 2-Methylbutane b.p. 28oC CH3CCH3 CH3 2,2-Dimethylpropane b.p. 9oC Individual van der Waals forces are very weak. However, a typical molecule can participate in so many polarization interactions that the van der Waals forces are among the most important of the intermolecular forces in the liquid phase. They are the only forces possible for nonpolar molecules. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 186 Daley & Daley Dipolar attractions are due to the uneven distribution of electrons in the covalent bonds that form between two atoms with differing electronegativities, thus creating a permanent dipole. The second category of attractions that occurs between molecules is dipolar attractions. Molecules with permanent dipoles have dipolar attractions because of the charge polarization in their bonds. The interactions between molecules with permanent dipoles are similar to the van der Waals interactions between molecules with induced dipoles. The only difference is that the dipoles are permanent. Methyl fluoride illustrates the interaction of molecules with a permanent dipole. Methyl fluoride has a very polar CF bond with a partial positive charge on the carbon and a partial negative charge on the fluorine atom: H HCF H In the liquid form, many other molecules of methyl fluoride surround each individual molecule of methyl fluoride. All these molecules tend to line up with the negative end of one dipole associated with the positive end of another: H HCF H H HCF H H HCF H H HCF H H HCF H Hydrogen bonding is an interaction between the hydrogen bonded to an electronegative atom (the hydrogen bond donor) and an atom with a nonbonding pair of electrons (the hydrogen bond acceptor). As with van der Waals forces, molecules with dipole attractions require energy to overcome these forces. The dipolar forces raise the boiling point of methyl fluoride above that of a comparable compound without electronegative substituents. For example, methyl fluoride and ethane, have similar molecular weights, but methyl fluoride boils at 78oC whereas ethane boils at 89oC. The third category of interactions that affects the boiling point is hydrogen bonding. Hydrogen bonding is a type of weak bonding interaction that involves a hydrogen bond donor and a hydrogen bond acceptor. A hydrogen bond donor is a molecule containing a hydrogen attached to an electronegative atom. The most common electronegative atoms in organic molecules are oxygen and nitrogen. A hydrogen acceptor bond is a molecule containing an atom with a nonbonding pair of electrons. The best hydrogen bond acceptors in organic molecules are also oxygen and nitrogen. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 187 Daley & Daley H O H H O H H O H H O H Hydrogen bonding in water The strongest hydrogen bonds are with the OH group. Weaker hydrogen bonds form with NH bonds. Much weaker still are the hydrogen bonds formed with SH and PH bonds. The strength of an individual hydrogen bond is roughly 5 kcal/mole, much smaller than the typical covalent bond strengths of 80-100 kcal/mole. Hydrogen bonds are stronger than dipolar interactions, which are about 1-2 kcal/mole. Of the three types of attractive forces, hydrogen bonding is the strongest. Hydrogen bonding substantially raises the boiling points of the compounds in which it occurs. For example, the isomeric compounds dimethyl ether and ethanol have widely different boiling points due to hydrogen bonding in ethanol. CH3 O CH3 CH3CH2 O H Dimethyl ether b.p. -23oC Ethanol b.p. 78oC Dimethyl ether has no hydrogen atoms attached to the oxygen, so no hydrogen bonding is possible. However, ethanol has a hydrogen attached to the oxygen, so hydrogen bonding occurs. Chemists consider hydrogen bonding a very weak or partial bonding between an oxygen of one molecule and a hydrogen of another. This bonding causes aggregations, or groupings, of molecules much like those resulting from dipolar attractions. However, these molecular aggregations possess much more stability than those resulting from dipolar interactions. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 188 Daley & Daley O H O H H O CH3 H O CH3 CH3 CH3 Hydrogen bonding in methanol Similar aggregations of molecules occur with amines. However, the boiling point differences of isomeric amines are less dramatic than for isomeric compounds of oxygen. For example, 3-methyl-1butanamine boils at 95-96oC whereas N,N-dimethylpropanamine boils at 65oC. NH2 3-Methyl-1-butanamine b.p. 95-96oC N N,N-Dimethylpropanamine b.p. 65oC The electronegativity of N is 3.04 and O is 3.44. Intramolecular interactions occur between different parts of the same molecule. The smaller difference in boiling points suggests that hydrogen bonds with NH bonds are weaker than hydrogen bonds with OH bonds. The NH bonds are less polar because nitrogen has a lower electronegativity than oxygen. The hydrogen bonds are weaker because the hydrogen end of the dipole in the NH bond is less positive than that in the OH bond. The previous discussion considers intermolecular hydrogen bonding. An additional factor comes into effect when two functional groups in one molecule participate in hydrogen bonding. The resultant intramolecular hydrogen bond is much more important than an intermolecular hydrogen bond in determining the properties of the molecule. For example, 2-nitrophenol has a much lower boiling point than either of its isomers, 3-nitrophenol or 4-nitrophenol. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 189 Daley & Daley O H O O H O O H N O N O N O O 2-Nitrophenol b.p. 215oC 3-Nitrophenol b.p. 263oC 4-Nitrophenol b.p. 279oC 2-Nitrophenol forms an intramolecular hydrogen bond between the hydrogen of the OH bond and one of the oxygens in the NO2 group. This intramolecular hydrogen bond prevents an intermolecular hydrogen bond from forming. Thus, boiling requires much less energy for the 2-nitrophenol isomer than for the 4-nitrophenol isomer because there are no strong intermolecular forces to overcome in going from the liquid phase to the gas phase. O H O Intramolecular hydrogen bond N O 2-Nitrophenol b.p. 215oC Section 3.5, page 000, discusses the relative stability of rings of various sizes. 2-Nitrophenol represents a category of compounds in which intramolecular hydrogen bonding forms either a stable five- or sixmembered ring. Section 3.5, page 000 discusses the relative stability of rings of various sizes. Because of this stability, they form more readily than rings of any other size. Any process that can result in a five- or six-membered ring is favored over one yielding another ring size. Expect to see this pattern repeatedly in your study of organic chemistry. Exercise 4.2 a) There is a 20oC difference in the boiling points of cyclohexane and 1,4-dioxacyclohexane (commonly called 1,4-dioxane). Which boils higher? Explain. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 190 Daley & Daley O Cyclohexane O 1,4-Dioxacyclohexane (1,4-Dioxane) b) Which has the higher boiling point, 3,3-dimethyl-1-butanol or 3,3-dimethyl-2-butanol? Explain. OH OH 3,3-Dimethyl-1-butanol 3,3-Dimethyl-2-butanol c) 1,6-Hexanediol has a boiling point of 250oC; 2,5-hexanediol has a boiling point that differs from this by 33oC. What is the boiling point of 2,5-hexanediol? Explain. OH HO OH OH 2,5-Hexanediol 1,6-Hexanediol Sample solution (a) Cyclohexane is a nonpolar compound. Therefore, the only intermolecular attractions that must be overcome for it to boil are van der Waals forces. 1,4-Dioxane has four polar CO bonds, so it participates in some dipolar attractions. Because of these added intermolecular attractions, 1,4-dioxane boils at a higher temperature. 4.4 Solubility Most chemical reactions take place in solution. That is, the solutes, or the reagents that you want to react, are uniformly mixed with, or dissolved in, the solvent. There are two types of solutions: single-phase solutions and multiple-phase solutions. In a singlephase, or homogeneous, solution all the solutes are soluble in the same solvent. Reactions take place best in single-phase solutions. This In a single-phase solution all solutes dissolved in the solution are soluble in the same solvent. In a multiple-phase solution there are two or more mutually insoluble solvents present. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 191 Daley & Daley One compound is soluble in another compound when both form one phase without either pure compound being separately visible. uniform distribution allows a higher rate of reaction between the reactants because they have a greater amount of contact with each other. Single-phase solutions also allow you to control the solution concentrations and reaction conditions. Multiple-phase mixtures are mixtures in which the individual solutes are not soluble in the same solvent, thus, you must use more than one solvent. Each solvent then forms a different layer in the mixture. In multiple-phase mixtures, the reactants have less contact with each other so that the rate of reaction is generally slower than in a single-phase solution. One compound is soluble in another as a result of the various intermolecular forces present in both compounds. These intermolecular forces are van der Waals forces, hydrogen bonding, and polarity. Molecular weight also plays a part in the solubility of a compound. The higher the molecular weight the lower the solubility. A compound dissolves in a solvent when the interactions between the compound and the solvent are either similar in strength or stronger than those interactions between molecules in the compound to be dissolved. There are also correlations between the molecular structure of the compounds and their solubility. A simple rule of thumb for estimating solubility is: A compound dissolves most easily in a solvent that is structurally similar to itself. The phrase structurally similar means that the solvent and the solute are similar types of molecules. For example, polar solutes or solvents mix with other polar solutes or solvents and nonpolar solutes or solvents mix with other nonpolar solutes or solvents, but polar solutes or solvents do not generally mix with nonpolar solutes or solvents. Water is highly polar making it an excellent solvent for ionic compounds and small polar organic molecules. If an ethyl group replaces one of the hydrogens in water, the result is ethanol, a solvent that has a substantially reduced polarity. Ethanol actually has both a polar and a nonpolar end. Although ethanol participates in hydrogen bonding and thus is a moderate solvent for salts and highly polar molecules, it also has van der Waals forces, making it a good solvent for a variety of organic molecules as well. Ethanol is too polar, however, to readily dissolve many low-polarity substances. Replacing both hydrogens in water with ethyl groups produces diethyl ether. Diethyl ether is only moderately polar. It acts as a hydrogen bond acceptor and so is a poor solvent for salts. However, it is a good solvent for a variety of both polar and nonpolar organic molecules. Water is a poor solvent for low polarity compounds, such as gasoline. Gasoline, a mixture of alkanes, will not dissolve in water. All attempts to dissolve gasoline in water, or water in gasoline, result in a two-phase solution with the gasoline floating on the water. Gasoline won't dissolve in water for several reasons. (1) The molecular structures of the two substances are different. Gasoline, a mixture of hydrocarbons, is nonpolar, and water is highly polar. (2) The www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 192 Daley & Daley molecular weights of the hydrocarbons are much larger in comparison to water. (3) The van der Waals forces between the molecules of gasoline are much stronger than they are in water. With water, primarily hydrogen bonding and secondarily dipolar attractions determine its intermolecular interactions. Both hydrogen bonding and dipolar attractions are stronger than van der Waals forces. Thus, gasoline does not dissolve in water because such a solution would reduce the number of stronger interactions. With gasoline and water, there is an unfavorable solvent-solute interaction. Solubility is dependent on the temperature of the solution. In general, the higher the temperature the higher the solubility of a solute in a given solvent. For example, the solubility of benzoic acid in water is 1.7 g/L at 0oC; it increases to 68.0 g/L at 95oC. Figure 4.7 shows a plot of the solubility of benzoic acid in water versus temperature. Note that the solubility of benzoic acid increases rapidly as the temperature increases above75oC. 80 70 60 50 40 30 20 10 0 0 20 40 60 80 10 0 Temperature Figure 4.7. A plot of the solubility of benzoic acid in water versus the temperature of the solution. Solubility Recrystallization A common laboratory purification technique is recrystallization. To begin, dissolve a solid compound in a minimum quantity of a hot solvent. On cooling, crystals of the original molecule form. The ideal solvent is one that does not dissolve the solute very well at low temperatures, but dissolves it readily at its boiling point. In general, the best solvent has a slightly lower polarity than the solute and thus the intermolecular interactions of solute with solvent are weaker than the intermolecular interactions of the solute. At higher temperature the solute becomes more soluble. Another useful characteristic for the solvent is a low boiling point, which makes it easier to remove solvent traces from the purified solute crystals. Finally, a good solvent should be a better solvent for any impurities than it is for the desired compound. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 193 Daley & Daley Table 4.2 summarizes the solubility trends for several common solvents and solutes. Solvent Diethyl ether (CH3CH2)2O Water Solute NaCl Ethanoic Acid 2-Propanone 2-Decanone 2-Propanol 2-Decanol Glycerol (1,2,3-Propanetriol) Chloromethane 1-Chloropropane Table 4.2. Solubility trends. H2O Ethanol CH3CH2OH Pentane CH3(CH2)3CH3 + + + + + s - s + + + + + + + + + + + + + s + + + + + + + s + + Symbols used: + = soluble, s = slightly soluble, - = insoluble Exercise 4.3 a) Is cyclohexane more or less soluble in water than 1,4dioxacyclohexane (more commonly called 1,4-dioxane)? Explain. O Cyclohexane O 1,4-Dioxacyclohexane (1,4-Dioxane) b) Which of the two following compounds is more soluble in pentane? (Hint: Be sure to draw a Lewis structure of each.) NH3Cl NH2 Cl Cyclohexyl ammonium chloride 4-Chlorocyclohexanamine c) Is methyl acetate more or less soluble in pentane than is propanoic acid? Explain. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 194 Daley & Daley O CH3COCH3 Methyl acetate O CH3CH2COH Propanoic acid Sample solution a) Cyclohexane is much less soluble in water than is 1,4-dioxane. The four CO bonds of 1,4-dioxane are available for dipolar attractions and act as hydrogen bond acceptors. [SIDEBAR] Surfactants When chemists hear the word surfactant, they first think of soaps or detergents. Although soaps and detergents are the most familiar of the surfactants, they are not the only ones. There are different kinds of surfactantsmany of which have commercial uses. Even your body produces surfactants. A surfactant is one of a class of chemical structures with a dual personality in respect to solubility. Surfactants are usually large molecules. One end is soluble in water; the other end is soluble in the typical organic solvents. This dual solubility is due to a highly polar segment at one end of the molecule and a nonpolar segment at the other end. Now look at soaps and detergents. A typical soap is the sodium or potassium salt of a long-chain carboxylic acid. Potassium stearate, CH3(CH2)16COOK, is a soap molecule. The head of the molecule, the ionic portion, is water-soluble; the tail of the molecule, the hydrocarbon chain, is water-insoluble. Chemists call the head the hydrophilic portion of the molecule and the tail the hydrophobic portion. Adding a soap molecule to water results in the formation of small agglomerations of molecules with long hydrophobic tails. These tails are dissolved in one another and the polar portions pointing outward into the water. These agglomerations are called micelles. Figure 4.8 illustrates the structure of a micelle. A surfactant is a molecule where one end of the molecule is soluble in one solvent and the other end of the same molecule is soluble in a different solvent. Surfactant is an acronym derived from the words surface active agent. Hydrophilic means water-loving. Hydrophobic means water-hating. A micelle is an agglomeration of molecules that contain polar and nonpolar portions with one set of portions dissolved in one another and the other set of portions dissolved in the solvent. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 195 Daley & Daley COO OOC COO COO - OOC COO OOC OOC COO COO - COO - OOC - OOC Figure 4.8. Soap molecules form a micelle in water. An emulsion is a suspension of small droplets of one liquid in another liquid in which it is normally insoluble. Hard water has any of the possible metal cations other than sodium or potassium ions dissolved in it. When you wash your skin or clothing, you want to remove two basic types of materials: dry particles and grease. Usually, you can readily rinse dry particles off. Washing grease off is much harder because it is insoluble in water. Here the surfactants in soap help. The tails of the soap molecules are a solvent for the grease. So they dissolve the grease. The heads of the soap molecules project out of the grease particle. Because the heads are soluble in water, they dissolve in the water, and the water carries the soap and the grease away. The interaction of the ionic ends of the surfactant molecules with the surrounding water holds the micelles suspended in the water (see Figure 4.8). The suspended micelles form an emulsion of the greasesoap solution and water. A difficulty with soaps is that they don't work well in hard water. The most common cations in hard water are calcium, magnesium, and iron ions. When you put soap into hard water, a precipitate of calcium, magnesium, or iron salts forms from an interaction of the cations with the carboxylate ions in the soap. This precipitate is the bathtub ring that you must scrub from your tub after you bathe using soap in hard water. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 196 Daley & Daley A wetting agent causes a solvent to adhere to, or wet, a surface better than it does when the solvent is pure. To solve this problem, chemists have developed a variety of surfactants with more soluble calcium or magnesium salts. One is sodium dodecanyl sulfate (sodium lauryl sulfate), CH3(CH2)11OSO3 Na. The textile industry uses sodium lauryl sulfate as a detergent. It is also used as a wetting agent in photography and toothpaste. In an aqueous solution, a wetting agent works by lowering the surface tension of the solution below that of pure water. A very important surfactant that lowers the surface tension of a solution is found in the transfer of oxygen to the bloodstream. This surfactant is necessary for life itself. A complex mixture of lipids and water coats the interior of the lungs. Simply described, this mixture consists of an aqueous solution of diapalmitoylphosphatidylcholine (DPPC), whose molecules have a hydrophilic head and a hydrophobic tail. O CH3(CH2)14 CH3(CH2)14 C C O O O CH2 CH CH2 O O P O Dipalmitoylphosphatidylcholine (DPPC) CH3 O CH2CH2N CH3 CH3 DPPC lowers the surface tension of the interior of the lungs, thereby increasing the rate of oxygen absorption. Polluted air and, particularly, cigarette smoke reduce the ability of the lungs to produce DPPC and thus inhibit the transfer of oxygen into the body. DPPC also keeps your lungs inflated. DPPC coats the inside of the alveoli in your lungs and this coating lowers the surface tension on the inside of the alveoli. Outside the alveoli is the blood, which has a much higher surface tension. This difference in surface tension pulls the alveoli into a spherical shape. Before birth, the fetus makes some respiratory-type movements, but its lungs remain collapsed. Immediately after birth the infant makes several strong inspiratory movements, and the lungs expand. The surfactant in the lining of the lungs keeps them inflated. However, when an infant is born before the surfactant system is functional, their lungs will not properly inflate or remain inflated. This problem, called hyaline membrane disease, is especially serious for premature infants who don't have a functional surfactant system. To help fight this condition, premature infants are given surfactants until their own surfactant system begins to function properly. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 197 Daley & Daley 4.5 Density Density is the mass of some substance per unit volume of that substance. An extraction is the transfer of a solute from one solvent to a better one. For liquids, density is usually measured as the number of grams per milliliterexpressed as g/mL. Density is a useful tool in the laboratory. Its most common use is in extractions, the transfer of a solute from one solvent to a better one. To do an extraction, choose a solvent that is better than the one the solute is currently dissolved in. The two solvents must be insoluble in one another. Then shake the two together and separate them. Because you know the density of each solvent and which solvent best dissolves the solute, you know which one to keep. You may also use density, as you do boiling or melting points, to help identify unknown liquids. Density of Solids You can easily estimate the relative density of many organic liquids, but you cannot do so as readily with solids because the density of a solid depends on too many factors. Two generalizations that relate to the density of solids are: organic solids typically have lower densities than inorganic solids, and nonionic solids typically have lower densities than ionic solids. Beyond these generalizations, this book does not consider the density of solids. Selected van der Waals radii, are listed in Table 3.1, page 000. The density of an organic liquid depends on three factors: the molecular weight of the substance, the ratio of the number of heavy atoms to the number of carbons in the molecule, and how well the molecules pack together. The first two factors are very closely related because, when determining the density of a liquid, there is more to the picture than just the molecule's actual weight. Of greater importance is the ratio of the number of heavy atoms to carbon atoms in the molecule. For example, the van der Waals radius of a methyl group is 200 pm, and the van der Waals radius of a bromine atom is 195 pm. The formula weight of a methyl group is 15, and the atomic weight of the bromine is 80. Bromine weighs more than five times as much as the methyl group, yet it takes up a smaller volume of space. This leads to a dramatic increase in density for bromine compared to a methyl group. Thus, when you compare the two in compounds pentane and 1bromobutane, for example, you find that both compounds have nearly the same molecular volume, but the density of pentane is 0.62 g/mL, and that of 1-bromobutane is 1.27 g/mL. This type of difference holds true for any set of molecules with similar sizesthe molecule with the heavier atoms is the molecule with the higher density. For example, cyclohexane and cyclohexaned12 (in which all of the hydrogens have been replaced with deuterium, 2H) have densities of 0.78 and 0.89 g/mL respectively. The third factor that determines a liquid's density is how efficiently its molecules pack together. The closeness of packing www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 198 Daley & Daley depends on how readily each molecule fits into an aggregation of molecules. For example, hexane, with a density of 0.66 g/mL, has many more stable conformations than does cyclohexane, with a density of 0.78 g/mL. Cyclohexane packs more efficiently because it is more symmetrical than hexane. Intermolecular attractions affect the packing efficiency of a molecule. The greater these attractions, the more readily the molecules pack together. Dipolar attractions and, especially, hydrogen bonding increase the density of a liquid. For example, compare the isomers diethyl ether and 1-butanol. Diethyl ether, which has van der Waals and dipolar attractions, has a density of 0.71 g/mL. 1-Butanol, which has hydrogen bonding in addition to van der Waals and dipolar attractions has a density of 0.81 g/mL. CH3CH2OCH2CH3 Diethyl ether 0.71 g/mL CH3CH2CH2CH2OH 1-Butanol 0.81 g/mL Chemists often perform extractions using a two-phase liquid system consisting of water and some organic liquid. They use an estimate of the density of the organic liquid relative to that of water. They work with estimates rather than precise densities because that is all they need to know to decide which layer is which. Here is a rule for estimating the densities of molecules relative to water. Any molecule containing bromine, iodine, or multiple chlorine atoms (e.g., CHCl3) has a density greater than that of water (1 g/mL). All other organic solvents generally have a density less than that of water. This rule, of course, is less effective for large molecules. In a large molecule, the functional group becomes less important, and all the physical properties tend to converge to the same values. Table 4.3 illustrates how chain length affects densities. Note that the density changes rapidly for molecules containing between two and ten carbons, but above ten carbons, the density does not change very much. Chain Length 10 14 0.730 0.763 0.783 0.818 0.870 0.866 0.830 0.834 0.865 0.857 Functional Group Alkane 1-Amino 1-Chloro 1-Phenyl Ethyl ester 2 0.509 0.683 0.898 0.783 0.900 6 0.660 0.766 0.879 0.814 0.871 18 0.789 0.822 0.864 Solid 0.852 Table 4.3. The effect of changing chain length on the density of liquids containing various functional groups. Exercise 4.4 www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 199 Daley & Daley Predict which compound of the following pairs has the higher density. CH 3 a) C H 3 CCH 2 CH 2 I or I CH 3 b) Bromotrifluoromethane or Tribromofluoromethane c) or CH3CH2OCH2CH3 O Sample solution b) Tribromofluoromethane has a much higher density than bromotrifluoromethane because it has three bromine atoms and bromotrifluoromethane has only one. Key Ideas from Chapter 4 The four phases of matter are solid, liquid, gas, and plasma. A sample compound in its gas phase has no definite shape or volume. In its liquid phase, it has volume but no specific shape. In its solid phase, it has both volume and shape. The molecular structure determines a compounds physical properties. Chemists are not yet able to correlate molecular structure with physical properties quantitatively. Intermolecular forces are the primary factors that determine the physical properties of a compound. These forces include van der Waals forces, dipolar interactions, and hydrogen bonding. The melting point of a compound is the temperature at which the compounds solid phase is in equilibrium with its liquid phase. The boiling point of a compound is the temperature at which its liquid phase is in equilibrium with its gas phase. Momentary dipolar attractions between adjacent nonpolar molecules cause van der Waals forces. www.ochem4free.com 5 July 2005 Organic Chemistry - Ch 4 200 Daley & Daley A carbon bonded to some atom of different electronegativity results in a polar bond. Polar bonds cause dipolar interactions. The interaction between a hydrogen that is attached to an electronegative atom and the nonbonding electrons of an atom in another molecule results in hydrogen bonding. If the intermolecular forces between solute and solvent are as strong or stronger as the attractions between molecules of the solute or the solvent itself, then the solute will be soluble in the solvent. The atomic weights of the atoms of a molecule contribute to the density of that molecule. When a molecule has a relatively large number of heavy atoms, its density is high. Another factor that determines the density of a compound is the strength of its intermolecular attractions. The stronger the intermolecular attractions, the higher the density. The larger the molecule, the less important are any functional groups in determining that molecules physical properties. www.ochem4free.com 5 July 2005 ... View Full Document

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