Effect of Molecular Architecture on Physical Properties:

Introduction

In the preceding chapters we discussed several of the significant features of bonding and the structures of organic molecules in detail. But have we thought about how we found the character and structure of a molecule? One respond to this question could be comparing its physical and chemical properties through those of the recognized compounds. Previous techniques of classification comprised the determination of physical properties these as melting point, boiling point, solubility and refractive index. The chemical techniques utilized for identification involved, though, either the degradation of the molecule to easy compounds of recognized structure or its synthesis from the simple compounds of known structure.

In this chapter, we will talk about the connection between molecular structure and physical properties. The study of physical properties is as well significant in the purification of organic compounds.

Molecular Architecture and Physical Properties:

The bonding and structural attributes of a compound are manifested in its physical properties. Therefore, physical properties of a compound these as melting point, boiling point, solubility, and so on, often provide expensive clues about its structure. Equally, if the structure of a compound is recognized, its physical properties can be predicated. The physical properties of a compound based upon the number and nature of atoms constituting its structural chapters and as well on the nature of forces holding such chapter jointly. We know that in case of ionic compounds, the positive and negative ions are held mutually via strong electrostatic forces. Contrary to this, in covalent compounds, the molecules are detained together via intermolecular forces. Let us now learn briefly what these intermolecular forces are. Then, we will learn how these intermolecular forces affect the physical properties of the compounds.

Intermolecular forces:

The three important intermolecular forces are: (i) dipole-dipole interactions, (ii) London forces and (iii) hydrogen bonding. Let us now consider these intermolecular forces one by one.

(i) Dipole-Dipole Interactions are defined as the interactions between the different molecules of a compound having permanent dipoles. Consider the example of chloromethane which has a permanent dipole. The molecule of chloromethane orient themselves in such a way that the positive end of 1 dipole points towards, and is therefore magnetized via, the negative end of the other dipole. These interactions, termed dipole-dipole interaction are depicted in Fig.

Fig: a) Dipole-dipole interactions between chloromethane molecules.

The dipole-dipole interactions are weak interactions and are of the order of 4 to 12 kJ mol^-1 whereas the bond energy for an ordinary covalent bond ranges from 125 to 420 kJ mol^-1

(ii) London forces: The intermolecular interactions exist between no polar molecules also. Consider two non polar molecules A and B in which the centre of positive charge coincides with that of the negative charge.

Whenever the molecules A and B approach each other, there is a distortion in the distribution of the charge resulting in a small and momentary dipole in one molecule. This small dipole can then create another dipole in the 2^nd molecule that is termed induced dipole. Therefore, if

the momentary dipole of molecule A is as exposed below;         

These a distribution of charge directs to mutual attraction between the molecules.

These induced dipole - induced dipole interaction are as well recognized as London forces. London forces are the only forces of attraction possible between nonpolar molecules. Such interactions are weaker than the dipole-dipole communications and are of the order of 4 kJ mol^-1. Such forces fluctuate through the distance between the molecules. If 'r' is the distance between the 2 molecules, then the London forces are proportional to 1/r6.

(iii) Hydrogen bonding consequences when a hydrogen atom is covalently connected to a strongly electronegative atom these as oxygen, nitrogen or fluorine. These hydrogen atoms have a huge affinity for the non bonded electrons of oxygen (or nitrogen or fluorine) atom of the other molecule. This kind of intermolecular interaction is recognized as hydrogen bonding. The hydrogen bonding is an individual kind of dipole-dipole interactions. The hydrogen bonding in the case of ethanol is symbolized below:

Note that the hydrogen bonds are indicated via the dash lines whereas the covalent bonds are symbolized by solid lines. The hydrogen bonding is a stronger interaction as evaluated to the dipole-dipole interactions but it is weaker than a covalent bond. The strength of hydrogen bond varieties from 10 - 40 kJ mol^-1. Hydrogen bonding has an important influence on physical properties such as melting point, boiling point and solubility of substances. This will be illustrated using examples in the following subsections.

The dipole-dipole, induced dipole-induced dipole etc. interactions are collectively known as van der Waals forces. Several authors desire to provide the name van der Waals forces only for London forces. Having understood the intermolecular forces, Allow us now learn how the variation in molecular structure influences such intermolecular forces that in turn in reflected in the physical properties of the molecules.

Melting point

The melting point of a material can be described as the temperature at which it undergoes the conversion from the solid to the liquid state. Pure crystalline solids contain spiky melting points. Therefore, melting point is utilized as a significant physical property together for the identification of organic compounds and for making the common assessment of the purity of such compounds.  Pure crystalline solids contain sharp melting points and they melt over a temperature range of 1^o or less. In contrast to this, contaminated crystalline solids melt over wider ranges of temperatures. In a crystalline solid, the constituent ions or molecules are positioned in an orderly and inflexible fashion. When these as solid are heated, the thermal energy of the molecules amplifies. This lastly leads to the disintegration of the crystal structure and at the melting point a unruly and random arrangement of particles, characteristic of a liquid, is attained. Because the electrostatic forces holding the ions are extremely strong, they can be defeat only at high temperatures. Hence, the ionic compounds usually have high melting points. For instance, the melting point of sodium chloride is 1074 K and that of sodium ethanoate is 595 K. But, the intermolecular forces are extremely weak as compared to the intrinsic forces and therefore, such can be overcome at lower temperatures leading to lower melting points for covalent compounds. The melting point of methane, a covalent compound, is only 90 K and the melting point of methanol, an additional covalent compound, is 179 K.

Let us now learn the consequence of molecular weight on the melting point. The association between molecular weight and melting point for alkanes is demonstrated in Fig.

Fig: Plot of melting points of straight chain alkanes alongside the molecular weight; the number of carbon atoms present in alkane molecule are as well indicated.

We can see in the figure that the melting point increases with the increase in the molecular weight. This can be explained due to increase in the London forces between the larger molecules of higher molecular weight. Thus, each additional methylene (-CH₂) unit contributes to the increase in melting point.

In a homologous series, the higher the molecular weight, the larger will be the molecules and the greater will be the 'area of contact' between the two molecules and hence the greater will be the London forces. You must have noticed in Fig. 1.2, the alternating pattern of melting points for the alkanes having odd and even number of carbon atoms. It is also evident from the figure that the compounds having even number of carbon atoms lie on a higher curve as compared to the compounds having odd number of carbon atoms. This can be explained on the basis that in solid state, the London forces among the molecules having odd number of carbon atoms are weaker than those in the molecules having even number of carbon atoms. This is because the molecules of alkanes having odd number of carbon atoms do not fit well in the crystal lattice as compare to those of the alkanes having even number of carbon atoms. After studying the effect of molecular weight on melting point, let us now see how the isomeric compounds having the same molecular weight, show different melting points. The melting points of straight chain and branched chain isomers of butane are given below

The branching of the carbon chain interferes through the regular packing of the molecules in the crystal; branched chain hydrocarbons tend to have lower melting points than their straight chain isomers.

But, in case, the branched molecule has a substantial equilibrium, then its melting point is moderately high. This is obviously evident whenever we evaluate the melting points of isomeric pentanes which are as specified below:

The branching from pentane to 2-methylbutane lowers the melting point but additional branching in 2, 2-dimethylpropane enhances the melting point. This can be explained through the reality that the symmetrical molecules fit jointly more simply in the crystal lattice and therefore have higher melting points as contrasted to the less symmetrical molecules. Therefore, higher melting point for 2, 2-dimethylpropane is justified. This is as well reflected when we analyse the melting points of cis - and trans-isomers. The trans - isomer being more balanced, fits improved in the crystal lattice than the less symmetrical cis - isomer. Therefore, the trans -isomers normally have higher melting points.

The nature of the functional collections present in a molecule as well influences its physical properties. For instance, whenever the functional group is these that it establishes polarity, and therefore leads to a permanent dipole instant in the molecule; then, due to the dipole-dipole forces of attraction between the polar molecules, they demonstrate higher melting points than the nonpolar molecules of similar molecular weights. For instance, the melting point of propanone, a polar molecule having molecular weight of 58, is 178 K. We can contrast it through the melting points of isomers of nonpolar butane (mol. Wt. = 58) we have just calculated above. This leads to the termination that the polar propanone has higher melting point than the nonpolar isomeric butanes.

The consequence of hydrogen bonding on melting point is little. But, the hydrogen bonding has important consequence on the boiling point, about which we will learn in the subsequent subsection.

Boiling point

The boiling point of a substance is the temperature at which it changes from the liquid to the gaseous state. At the boiling point the vapour pressure of a liquid is equivalent to the external pressure. Therefore, the boiling point based on the external pressure and it enhances through amplify in the external pressure. Therefore, while reporting the boiling point of a material, external pressure must be particular. Generally, the boiling points are reported at full of atmosphere pressure. Analogous to the case of melting points, the boiling points are as well utilized as stables for identification and classification of liquid materials. The information of boiling points is as well significant in the purification of liquids. Let us now learn several of the causes affecting the boiling point. The boiling point of a material based on its molecular structure. In a homologous sequence, the boiling points of the compounds enhance through the amplify in the number of carbon atoms. In other terms, we can say that the boiling point increase with increase in molecular weight. Generally, this enhance in boiling point amounts to 20-30^o for the addition of each carbon atom in the molecule. The enhance in boiling point through molecular weight can be again traits to enhanced London forces of attraction between superior molecules.

Among isomeric molecules, because the unbranched isomer is linear and therefore enlarged in shape, it has better surface area as compared to the branched isomers. Hence, the London forces are stronger in the unbranched isomer leading to higher boiling point for this isomer. This is demonstrated in Fig for the isomers of butane.

Fig: A comparison of intermolecular interactions for straight chain and branched chain isomers of butane

The polarity of a compound as well influences its boiling point. When we compare molecules having similar shape and size, the more polar molecule has the higher boiling point. Instances are:

Alcohols have uncommonly high boiling points as contrasted to the other compounds of analogous molecular weight or size. For instance, ethanol CH₃CH₂OH, which has similar molecular formula as that of dimethyl ether, has the boiling point 351 K. This can be described due to hydrogen bonding. Hydrogen bonding for ethanol has been demonstrated previous in sub-sec. Therefore, to vaporize these as compound; hydrogen bonds between the molecules must be broken. This needs energy that is manifested as the uncommonly high boiling point for these compounds.

The hydrogen bonding as given for ethanol is recognized as intermolecular hydrogen bonding that means that the hydrogen bonds are present between the molecules.

Hydrogen bonding can as well happen within similar molecule in which case it is termed intermolecular hydrogen bonding. Therefore, 2-hydroxybenzaldehyde can structure only intermolecular hydrogen bonds. The enhanced intermolecular attraction due to intermolecular hydrogen bonding is reflected in the superior boiling point for 4-hydroxybenzaldehyde as contrasted to 2-hydroxybenzaldehyde in which this intermolecular interaction is absent. Hydrogen bonding is as well significant in other ways. As we shall see in the next part, hydrogen bonding plays a significant role in the solubility of organic compounds.

Solubility

When any material dissolves in a solvent, its constituent ions or molecules get divided from each other and the space between them is filled via solvent molecules. This is identified a solvation and the amount of material dissolved in a certain amount of solvent is termed to as its solubility in that solvent. Solubility therefore based on the interactions between solute-solute, solute-solvent and solvent-solvent molecules. Obviously strong solute-solvent molecular interactions as compare to those of solute - solute or solvent - solvent molecules will lead to dissolution of the solute. Alike to the procedures of melting or boiling, dissolution n of a substance as well needs that the interionic or intermolecular forces of attraction between the ions or molecules must be defeat. The strong electrostatic forces between the ions of an ionic compound can be overcome through the solvents which have high dielectric constant. Therefore, water which has a high dielectric constant of 80, dissolves ionic compounds readily whereas solvents as carbon tetrachloride (? =1.2) or ether (? = 4.4) are very poor solvents for these compounds. Therefore ionic compounds have superior solubility in polar solvents.

The dielectric constant ?, of a solvent determines its ability to divide the ions of the solute. The word polar has double usage in organic chemistry. When we refer that it has an important dipole moment, µ. But, when we discuss a polar solvent, we comprehend that it has a high dielectric constant, ?. Therefore, the dipole instant is the property of individual molecules whereas solvent polarity or dielectric constant is a property of many molecules performing jointly.

In finding out the solubility of covalent compounds, the rule of thumb is like-dissolves-like. Because water is a polar compound, it is a superior solvent for polar compounds, but is a poor solvent for hydrocarbon which is nonpolar in nature. Therefore, the hydrocarbons readily dissolve in additional hydrocarbons or in nonpolar solvents these as benzene, ether or tetrahydrofuran. As most organic compounds have both a polar and a nonpolar part, their solubility will depend upon the balance between the two parts. Consider the solubilities of three alcohols, ethanol, butanol and hezanol in water, as given below:

 Nonpolar              part polar part                      nonpolar part                          nonpolar part

Solubility: Miscible through water 3.9 g in 1 dm³ of water 5.9 g in 1 dm³ of water in all proportions.

We can observe that as the size of the nonpolar portion of the molecule enhances its solubility in water reduces.

The solubility of organic compounds in water as well based on the extent of hydrogen bonding possible between the solute and the solvent (water) molecules.

For instance, the greater solubility f ether in water as compared to that of pentane (in water can be accounted on the basis of hydrogen bonding present in the former case.

Because the olefinic, acetylenic or benzenoid character doesn't influence the polarity much, the solubility of unsaturated and aromatic hydrocarbons in water is alike to that of alkanes. In compounds as ethers, esters, aldehydes, ketones, alcohols, amides, acids and amines, solubility in water based on the length of the alkyl chain and the members enclosing less than 5 carbon atoms in the molecules are soluble in water.

Enhance in the intermolecular forces in a solute, as a consequence of amplify in the molecular weight, is as well reflected in the low solubility of compounds having high molecular weight. For instance, glucose is soluble in water but its polymer, starch is insoluble in water Polymers have elevated molecular weight. Therefore, in a homologous series, the solubility of the members reduces through the amplify in molecular weight.

Though, branching of the carbon chain leads to a decrease in the intermolecular forces. Hence, the branched chain isomer is more soluble as compared to the straight chain isomer. Apart from other factors discussed above, solubility of a compound in a given solvent usually raises through temperature. Sometimes high solubility of a compound is examined due to a chemical reaction which performs as a driving force. One these groups of reactions are acid-base reactions. For instance, the higher solubility of aniline in aqueous acid is due to the formation of anilinium ion.

Even though determination of the physical properties these as those discussed above aids in the identification of organic compounds, physical techniques involving the utilize of spectroscopy permit determination of the molecular structure much more quickly and nondestructively using little quantities of material.

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