Predicting Bond Type From Electronegativities

Intermolecular forces are electrostatic in nature. Molecules are composed of electrically charged particles (electrons and protons), and it is common for them to have regions that are predominantly charged positive or negative. Attractive forces between molecules arise when electrostatic forces attract positive regions on one molecule to negative regions on another. The strength of the attractions between molecules depends on the polarities of chemical bonds within the molecules and the geometrical shapes of the molecules.

Chemical bonds — ionic, nonpolar covalent, and polar covalent: At the simplest level, the chemical bonds that hold atoms together in a molecule are of two types:

1. Ionic bonds: occur when one atom attracts an electron away from another atom to form a positive and a negative ion. The ions are then bound together by electrostatic attraction. The electron transfer occurs because the electron-receiving atom has a much stronger attraction for electrons in its vicinity than does the electron-losing atom.

2. Covalent bonds: are formed when two atoms share electrons, called bonding electrons, in the space between their nuclei. The electron-attracting properties of covalent bonded atoms are not different enough to allow one atom to pull an electron entirely away from the other. However, unless both atoms attract bonding electrons equally, the average position of the bonding electrons will be closer to one of the atoms. The atoms are held together because their positive nuclei are attracted to the negative charge of the shared electrons in the space between them.

When two covalent bonded atoms are identical, as in Cl2, the bonding electrons are always equally attracted to each atom and the electron charge is uniformly distributed between the atoms. Such a bond is called a nonpolar covalent bond, meaning that it has no polarity, i.e., no regions with net positive or negative charge.

When two covalent bonded atoms are of different kinds, as in HCl, one atom may attract the bonding electrons more strongly than the other. This results in a non-uniform distribution of electron charge between the atoms where one end of the bond is more negative than the other, resulting in a polar bond.

Figure 2.2 illustrates the electron distributions in nonpolar and polar covalent bonds. The strength with which an atom attracts bonding electrons to itself is indicated by a quantity called electronegativity. Electronegativities of the elements, shown in Table 2.1, are relative numbers with an arbitrary maximum value of 4.0 for fluorine, the most electronegative element. Electronegativity values are approximate, to be used primarily for predicting the relative polarities of covalent bonds.

The electronegativity difference between two atoms indicates what kind of bond they will form. The greater the difference in electronegativities of bonded atoms, the more strongly are the bonding electrons attracted to the more electronegative atom, and the more polar is the bond. The following "rules of thumb" usually apply, with very few exceptions.

Because electronegativity differences can vary continuously between zero and four, bond character also can vary continuously between nonpolar covalent and ionic, as illustrated in Figure 2.3.

Rules of Thumb (Use Table 2.1)

1. If the electronegativity difference between two bonded atoms is zero, they will form a nonpolar covalent bond. Examples are O2, H2, N2, and NCl.

2. If the electronegativity difference between two atoms is between zero and 1.7, they will form a polar covalent bond. Examples are HCl, NO, and CO.

3. If the electronegativity difference between two atoms is greater than 1.7, they will form an ionic bond. Examples are NaCl, HF, and KBr.

4. Relative electronegativities of the elements can be predicted by an element's position in the Periodic Table. Ignoring the noble gases:

a. The most electronegative element (F) is at the upper right corner of the Periodic Table.

b. The least electronegative element (Fr) is at the lower left corner of the Periodic Table.

c. In general, electronegativities increase diagonally up and to the right in the Periodic Table. Within a given Period (or row), electronegativities tend to increase in going from left to right; within a given Group (or column), electronegativities tend to increase in going from bottom to top.

d. The farther apart two elements are in the Periodic Table the more different are their electronegativities, and the more polar will be a bond between them.

H >CI

Nonpolar covalent bond established by a uniform distribution of bonding electrons between identical atoms in CI2 .

Polar bond established by a non-uniform distribution of bonding electrons between different atoms in HCI. Electron charge is concentrated toward the more electronegative atom, indicated by 8- The less electronegative atom is indicated by5+.

CI-CI

hi——CI

Nonpolar covalent bond indicated by a straight line joining the atoms.

Polar bond indicated by an arrow over the bond pointing to the more electronegative atom. A vertical cross line shows the more positive end of the bond. The length of the arrow indicates the magnitude of the bond's dipole moment.

FIGURE 2.2 Uniform and non-uniform electron distributions, resulting in nonpolar and polar covalent chemical bonds. The use of a delta (5) in front of the + and - signs signifies that the charges are partial, arising from a non-uniform electron charge distribution rather than from the transfer of a complete electron.

Dipole Moments

For polar bonds, we can define a quantity, called the dipole moment, which serves as a measure of the non-uniform charge separation. Hence, the dipole moment measures the degree of the bond polarity. The more polar the bond, the larger is its dipole moment. The dipole moment, m, is equal to the magnitude of positive and negative charges at each end of the dipole multiplied by the distance, d, between the charges.

Polarity arrows, as shown in Figure 2.4, are vector quantities. They show both the magnitude and direction of the bond dipole moment. The length of the arrow indicates how large is the dipole moment, and the direction of the arrow points to the charge separation.

60898 Curve
FIGURE 2.3 Bond character as a function of the electronegativity difference.

TABLE 2.1

Electronegativity Values of the Elements

TABLE 2.1

Electronegativity Values of the Elements

H

2

13

14

15

16

17

2.1

2A

3A

4A

5A

6A

7A

3

4

5

6

7

8

9

Li

Be

B

C

N

O

F

1.0

1.5

2.0

2.5

3.0

3.5

4.0

11

12

13

14

15

16

17

Na

Mg

3

4

5

6

7

8

9

10

11

12

Al

Si

P

S

Cl

1.0

1.2

3B

4B

5B

6B

7B

8B

8B

8B

1B

2B

1.5

1.8

2.1

2.5

3.0

19

20

21

22

23

24

25

26

27

28

29

30

31

32

33

34

35

K

Ca

Sc

Ti

V

Cr

Mn

Fe

Co

Ni

Cu

Zn

Ga

Ge

As

Se

Br

0.9

1.0

1.3

1.4

1.5

1.6

1.6

1.7

1.7

1.8

1.8

1.6

1.7

1.9

2.1

2.4

2.8

37

38

39

40

41

42

43

44

45

46

47

48

49

50

51

52

53

Rb

Sr

Y

Zr

Nb

Mo

Tc

Ru

Rh

Pd

Ag

Cd

In

Sn

Sb

Te

1

0.9

1.0

1.2

1.3

1.5

1.6

1.7

1.8

1.8

1.8

1.6

1.6

1.6

1.8

1.9

2.1

2.5

55

56

57

72

73

74

75

76

77

78

79

80

81

82

83

84

85

Cs

Ba

La

Hf

Ta

w

Re

Os

lr

Pt

Au

Hg

TI

Pb

Bi

Po

At

0.8

1.0

1.1

1.3

1.4

1.5

1.7

1.9

1.9

1.8

1.9

1.7

1.6

1.7

1.8

1.9

2.1

87

88

89

104

105

106

107

108

109

Fr

Ra

#Ac

Rf

Db

Sg

Bh

Hs

Mt

0.8

1.0

1.1

?

?

?

?

?

?

# Actinide series

*Lanthanide series

58

59

60

61

62

63

64

65

66

67

68

69

70

71

Ce

Pr

Nd

Pm

Sm

Eu

Gd

Tb

Dy

Ho

Er

Tm

Yb

Lu

1.1

1.1

1.1

1.1

1.2

1.1

1.2

1.1

1.2

1.2

1.2

1.3

1.0

90

91

92

93

94

95

96

97

98

99

100

101

102

103

Th

Pa

u

Np

Pu

Am

Cm

Bk

Cf

Es

Fm

Md

No

Lr

1.3

1.5

1.5

1.3

1.3

1.3

1.3

1.3

1.3

1.3

1.3

1.3

1.3

1.5

polarity arrow x

FIGURE 2.4 Molecular dipole moment as indicated by a polarity arrow.

2.9 MOLECULAR GEOMETRY, MOLECULAR POLARITY, AND INTERMOLECULAR FORCES

Knowing whether a molecule is polar or not helps to predict its water solubility and other properties. The presence of polar bonds in a molecule may make the molecule polar also. A molecule is polar if the polarity vectors of all its bonds add up to give a net polarity vector to the molecule. Like polar bonds, a polar molecule has a negatively charged region where electron density is concentrated, and a positively charged region where electron density is diminished. The polarity of a molecule is the vector sum of all its bond polarity vectors. A polar molecule can be experimentally detected by observing whether an electric field exerts a force on it that makes it align its charged regions in the direction of the field. Polar molecules will point their negative ends toward the positive source of the field, and their positive ends toward the negative source. To predict if a molecule is polar, we need to answer two questions:

1. Does the molecule contain polar bonds? If it does, then it might be polar; if it doesn't, it cannot be polar.

2. If the molecule contains polar bonds, do all the bond polarity vectors add to give a resultant molecular polarity? If the molecule is symmetrical in a way that the bond polarity vectors add to zero, then the molecule is nonpolar although it contains polar bonds. If the molecule is asymmetrical and the bond polarity vectors add to give a resultant polarity vector, the resultant vector indicates the molecular polarity.

Examples of Nonpolar Molecules

Nonpolar molecules invariably have low water solubility. A molecule with no polar bonds cannot be a polar molecule. Thus, all diatomic molecules where both atoms are the same, such as H2, O2, N2, and Cl2, are nonpolar because there is no electronegativity difference across the bond. On the other hand, a molecule with polar bonds whose dipole moments add to zero because of molecular symmetry is not a polar molecule. Carbon dioxide, carbon tetrachloride, hexachlorobenzene, para-dichloroben-zene, and boron tribromide are all symmetrical and nonpolar, although all contain polar bonds.

Carbon dioxide: Oxygen is more electronegative (EN(O2) = 3.5) than carbon (EN(C) = 2.5). Each bond is polar, with the oxygen atom at the negative end of the dipole. Because CO2 is linear with carbon in the center, the polarity vectors cancel each other and CO2 is nonpolar.

Carbon tetrachloride: EN(C) = 2.5, EN(Cl) = 3.0 C -+-> Cl.

Although each bond is polar, the tetrahedral symmetry of the molecule results in no net dipole moment so that CCl4 is nonpolar.

Cl CI Hexachlorobenzene: The bond polarities are the same as in CC14 above.

C6C16 is planar with hexagonal symmetry. All the bond polarities cancel one another and the molecule is nonpolar.

Ci h Para-dichlorobenzene: This molecule also is planar. It has polar bonds of

_two magnitudes, the smaller polarity H-h^C bond and the larger polarity

C-h^Cl bond. The H and CI atoms are positioned so that all polarity vectors cancel and the molecule is nonpolar. Check the electronegativity values in

Boron tribromide: EN(B) = 2.0, EN(Br) = 2.8 B-h^Br. BBr3 has trigonal planar symmetry, with 120° between adjacent bonds. All the polarity vectors cancel and the molecule is nonpolar.

Examples of Polar Molecules

Polar molecules are generally more water-soluble than nonpolar molecules of similar molecular weight. Any molecule with polar bonds whose dipole moments do not add to zero is a polar molecule. Carbon monoxide, carbon trichloride, pentachlorobenzene, ortho-dichlorobenzene, boron dibromochloride, and water are all polar.

Carbon monoxide: Oxygen is more electronegative (EN(O2) = 3.5) than carbon (EN(C) = 2.5). Every diatomic molecule with a polar bond must be a polar molecule.

Carbon trichloride: EN(C) = 2.5, EN(Cl) = 3.0, EN(H) = 2.1. It has polar bonds of two magnitudes, the smaller polarity H +—> C bond and the larger polarity C +—> Cl bond. The asymmetry of the molecule results in a net dipole moment, so that CHCl3 is polar.

Pentachlorobenzene: The bond polarities are the same as in CHCl3 above. The bond polarities do not cancel one another and the molecule is polar.

Ortho-dichlorobenzene: This molecule is planar and has two kinds of polar H bonds: H-h^C and C-h^Cl. The bond polarity vectors do not cancel, making the molecule polar.

q Boron dibromochloride: EN(B) = 2.0, EN(Br) = 2.8, EN(C1) = 3.0. In

| BBr2Cl, the polarity vectors of the polar bonds, B-h^Br and B-i—>C1, do not quite cancel and the molecule is slightly polar.

Br Br

Water: is a particularly important polar molecule. Its bond polarity vectors add to give the water molecule a high polarity (i.e., dipole moment). The dipole-dipole forces between water molecules are greatly strengthened by hydrogen bonding (see discussion below), which contributes to many of water's unique characteristics, such as relatively high boiling point and viscosity, low vapor pressure, and high heat capacity.

The Nature of Intermolecular Attractions

All molecules are attracted to one another because of electrostatic forces. Polar molecules are attracted to one another because the negative end of one molecule is attracted to the positive ends of other molecules, and vice versa. Attractions between polar molecules are called dipole-dipole forces. Similarly, positive ions are attracted to negative ions. Attractions between ions are called ion-ion forces. If ions and polar molecules are present together, as when sodium chloride is dissolved in water, there can be ion-dipole forces, where positive and negative ions (e.g., Na+ and Cl ) are attracted to the oppositely charged ends of polar molecules (e.g., H2O).

However, nonpolar molecules also are attracted to one another although they do not have permanent charges or dipole moments. Evidence of attractions between nonpolar molecules is demonstrated by the fact that nonpolar gases such as methane (CH4), oxygen (O2), nitrogen (N2), ethane (CH3CH3), and carbon tetrachloride (CCl4) condense to liquids and solids when the temperature is lowered sufficiently. Knowing that positive and negative charges attract one another makes it easy to understand the existence of attractive forces among polar molecules and ions. But how can the attractions among nonpolar molecules be explained?

In nonpolar molecules, the valence electrons are distributed about the nuclei so that, on average, there is no net dipole moment. However, molecules are in constant motion, often colliding and approaching one another closely. When two molecules approach closely, their electron clouds interact by electrostatically repelling one another. These repulsive forces momentarily distort the electron distributions within the molecules and create transitory dipole moments in molecules that would be nonpolar if isolated from neighbors. A transitory dipole moment in one molecule induces electron charge distortions and transitory dipole moments in all nearby molecules. At any instant in an assemblage of molecules, nearly every molecule will have a non-uniform charge distribution and an instantaneous dipole moment. An instant later, these dipole moments would have changed direction or disappeared so that, averaged over time, nonpolar molecules have no net dipole moment. However, the effect of these transitory dipole moments is to create a net attraction among nonpolar molecules. Attractions between nonpolar molecules are called dispersion forces or London forces (after Professor Fritz London who gave a theoretical explanation for them in 1928).

Hydrogen bonding: An especially strong type of dipole-dipole attraction, called hydrogen bonding, occurs among molecules containing a hydrogen atom covalently bonded to a small, highly electronegative atom that contains at least one valence shell nonbonding electron pair. An examination of Table 2.1 shows that fluorine, oxygen, and nitrogen are the smallest and most electronegative elements that contain nonbonding valence electron pairs. Although chlorine and sulfur have similarly high electronegativities and contain nonbonding valence electron pairs, they are too large to consistently form hydrogen bonds (H-bonds). Because hydrogen bonds are both strong and common, they influence many substances in important ways.

Hydrogen bonds are very strong (10 to 40 kJ/mole) compared to other dipole-dipole forces (from less than 1 to 5 kJ/mole). The hydrogen atom's very small size makes hydrogen bonding so uniquely strong. Hydrogen has only one electron. When hydrogen is covalently bonded to a small, highly electronegative atom, the shift of bonding electrons toward the more electronegative atom leaves the hydrogen nucleus nearly bare. With no inner core electrons to shield it, the partially positive hydrogen can approach very closely to a nonbonding electron pair on nearby small polar molecules. The very close approach results in stronger attractions than with other dipole-dipole forces.

Because of the strong intermolecular attractions, hydrogen bonds have a strong effect on the properties of the substances in which they occur. Compared with nonhydrogen bonded compounds of similar size, hydrogen bonded substances have relatively high boiling and melting points, low volatilities, high heats of vaporization, and high specific heats. Molecules that can H-bond with water are highly soluble in water; thus, all the substances in Figure 2.5 are water-soluble.

Comparative Strengths of Intermolecular Attractions

The strength of dipole-dipole forces depends on the magnitude of the dipole moments. The strength of ion-ion forces depends on the magnitude of the ionic charges. The strength of dispersion forces depends on the polarizability of the nonpolar molecules. Polarizability is a measure of how easily the electron distribution can be distorted by an electric field — that is, how easily a dipole moment can be induced in an atom or a molecule. Large atoms and molecules have more electrons and larger electron clouds than small ones. In large atoms and molecules, the outer shell electrons are farther from the nuclei and, consequently, are more loosely bound. The electron distributions can be more easily distorted by external charges. In small atoms and molecules, the outer electrons are closer to the nuclei and are more tightly held. Electron charge distributions in small atoms and molecules are less easily distorted.

Therefore, large atoms and molecules are more polarizable than small ones. Since atomic and molecular sizes are closely related to atomic and molecular weights, we can generalize that polar-izability increases with increasing atomic and molecular weights. The greater the polarizability of atoms and molecules, the stronger are the intermolecular dispersion forces between them. Molecular shape also affects polarizability. Elongated molecules are more polarizable than compact molecules. Thus, a linear alkane is more polarizable than a branched alkane of the same molecular weight.

All atoms and molecules have some degree of polarizability. Therefore, all atoms and molecules experience attractive dispersion forces, whether or not they also have dipole moments, ionic charges, or can hydrogen-bond. Small polar molecules are dominated by dipole-dipole forces since the contribution to attractions from dispersion forces is small. However, dispersion forces may dominate in very large polar molecules.

Rules of Thumb

1. The higher the atomic or molecular weights of nonpolar molecules, the stronger are the attractive dispersion forces between them.

2. For different nonpolar molecules with the same molecular weight, molecules with a linear shape have stronger attractive dispersion forces than do branched, more compact molecules.

3. For polar and nonpolar molecules alike, the stronger the attractive forces, the higher the boiling point and freezing point, and the lower the volatility of the substance.

(a) Water; extensive H-bonding gives water its high boiling point. When water freezes, H-bonding forces the molecules into an open solid structure, with the result that the solid form is less dense than the liquid. Thus, ice floats on water.

(b) Ammonia dissolved in water.

(c) Ethanol; hydrogens bonded to carbons, as seen in (c), (d) and (e), cannot form H-bonds because carbon is not electronegative enough.

(d) Ethanol dissolved in water.

(e) Acetic acid; pure acetic acid contains a high percentage of dimers (double molecules) held together by H-bonds between the -COOH groups.

(f) Hydrogen fluoride forms zigzag chains.

FIGURE 2.5 Examples of hydrogen bonding among different molecules.

Examples

1. Consider the halogen gases fluorine (F2, MW = 38), chlorine (Cl2, MW = 71), bromine (Br2, MW = 160), and iodine (I2, MW = 254). All are nonpolar, with progressively greater molecular weights and correspondingly stronger attractive dispersion forces as you go from F2 to I2. Accordingly, their boiling and melting points increase with their molecular weights. At room temperature, F2 is a gas (bp = -188° C), Cl2 is also a gas but with a higher boiling point (bp = -34° C), Br2 is a liquid (bp = 58.8° C), and I2 is a solid (mp = 184° C).

TABLE 2.2

Some Properties of the First Twelve Straight-Chain Alkanes Molecular Melting Pointa Boiling Point

TABLE 2.2

Some Properties of the First Twelve Straight-Chain Alkanes Molecular Melting Pointa Boiling Point

Alkane

Formula

Weight

oC

oC

methane

CH4

16

-183

-162

ethane

C2H6

30

-172

-89

propane

C3H8

44

-188

-42

n-butane

C4H10

58

-138

0

n-pentane

C5H12

72

-130

36

n-hexane

C6H14

86

-95

69

n-heptane

C7H16

100

-91

98

n-octane

C8H18

114

-57

126

n-nonane

C9H20

128

-51

151

n-decane

C10H22

142

-29

174

n-dodecane

C12H26

170

-10

216

a Deviations from the general trend in melting points occur because melting points for the smallest alkanes are more strongly influenced by differences in crystal structure and lattice energy of the solid.

a Deviations from the general trend in melting points occur because melting points for the smallest alkanes are more strongly influenced by differences in crystal structure and lattice energy of the solid.

2. Alkanes are compounds of carbon and hydrogen only. Although C—H bonds are slightly polar (electronegativity of C = 2.5; electronegativity of H = 2.1) all alkanes are nonpolar because of their bond geometry. In the straight-chain alkanes (called normal-alkanes), as the alkane carbon chain becomes longer, the molecular weights and, consequently, the attractive dispersion forces become greater. Consequently, melting points and boiling points become progressively higher. The physical properties of the normal-alkanes in Table 2.2 reflect this trend.

3. Normal-butane [n-C5H12] and dimethylpropane [CH3C(CH3)2CH3] are both nonpolar and have the same molecular weights (MW = 72). However, n-C5H12 is a straight-chain alkane while CH3C(CH3)2CH3 is branched. Thus, n-C5H12 has stronger dispersion attractive forces than CH3C(CH3)2CH3 and a correspondingly higher boiling point.

CH2 CH2

normal-pentane: bp = 36o C

Dimethylpropane: bp = 9.5oC

2.10 SOLUBILITY AND INTERMOLECULAR ATTRACTIONS

In liquids and gases, the molecules are in constant, random, thermal motion, colliding and intermingling with one another. Even in solids, the molecules are in constant, although more limited, motion. If different kinds of molecules are present, random movement tends to mix them uniformly. If there were no other considerations, random motion would cause all substances to dissolve completely into one another. Gases and liquids would dissolve more quickly and solids more slowly.

However, intermolecular attractions must also be considered. Strong attractions between molecules tend to hold them together. Consider two different substances A and B, where A molecules are attracted strongly to other A molecules, B molecules are attracted strongly to other B molecules, but A and B molecules are attracted weakly to one another. Then, A and B molecules tend to stay separated from each other. A molecules try to stay together and B molecules try to stay together, each excluding entry from the other. In this case, A and B are not soluble in one another.

As an example of this situation, let A be a nonpolar, straight-chain liquid hydrocarbon such as n-octane (C8H18) and let B be water (H2O). Octane molecules are attracted to one another by strong dispersion forces, and water molecules are attracted strongly to one another by dipole-dipole forces and H-bonding. Dispersion attractions are weak between the small water molecules. Because the small water molecules have low polarizability, octane cannot induce a strong dispersion force attraction to water. Because octane is nonpolar, there are no dipole-dipole attractions to water. When water and octane are placed in the same container, they remain separate forming two layers with the less dense octane floating on top of the water.

However, if there were strong attractive forces between A and B molecules, it would help them to mix. The solubility of one substance (the solute) in another (the solvent) depends mostly on intermolecular forces and, to a much lesser extent, on conditions such as temperature and pressure. Substances are more soluble in one another when intermolecular attractions between solute and solvent are similar in magnitude to the intermolecular attractions between the pure substances. This principle is the origin of the rules of thumb that say "like dissolves like" or "oil and water don't mix." "Like" molecules have similar polar properties and, consequently, similar intermolecular attractions. Oil and water do not mix because water molecules are attracted strongly to one another, and oil molecules are attracted strongly to one another; but water molecules and oil molecules are attracted only weakly to one another.

Rules of Thumb

1. The more symmetrical the structure of a molecule containing polar bonds, the less polar and the less soluble it is in water.

2. Molecules with OH, NO, or NH groups can form hydrogen bonds to water molecules. They are the most water-soluble non-ionic compounds, even if they are nonpolar because of geometrical symmetry.

3. The next most water-soluble compounds contain O, N, and F atoms. All have high electronegativities and allow water molecules to H-bond with them.

4. Charged regions in ionic compounds (like sodium chloride) are attracted to polar water molecules. This makes them more soluble.

5. Most compounds in oil and gasoline mixtures are nonpolar. They are attracted to water very weakly and have very low solubilities.

6. All molecules, including nonpolar molecules, are attracted to one another by dispersion forces. The larger the molecule the stronger the dispersion force.

7. Nonpolar molecules, large or small, have low solubilities in water because the small-sized water molecules have weak dispersion forces, and nonpolar molecules have no dipole moments. Thus, there are neither dispersion nor polar attractions to encourage solubility.

Examples

1. Alcohols of low molecular weight are very soluble in water because of hydrogen bonding. However, their solubilities decrease as the number of carbons increase. The -OH group on alcohols is hydrophilic (attracted to water), while the hydrocarbon part is hydrophobic (repelled from water). If the hydrocarbon part of an alcohol is large enough, the hydro-phobic behavior overcomes the hydrophilic behavior of the -OH group and the alcohol has low solubility. Solubilities for alcohols with increasingly larger hydrocarbon chains are given in Table 2.3.

TABLE 2.3

Solubilities and Boiling Points of Some Straight Chain Alcohols

TABLE 2.3

Solubilities and Boiling Points of Some Straight Chain Alcohols

Molecular

Melting Pointa

Boiling Point

Aqueous solubility

Name

Formula

Weight

(0C)

(° C)

at 25°C (mol/L)

Methanol

CH3OH

32

-98

65

<*> (miscible)

Ethanol

C2H5OH

46

-130

78

<» (miscible)

1-propanol

C3H7OH

60

-127

97

<» (miscible)

1-butanol

C4H9OH

74

-90

117

0.95

1-pentanol

C5H11OH

88

-79

138

0.25

1,5-pentanediolb

C5H10(OH)2

104

-18

239

<» (miscible)

1-hexanol

C6H13OH

102

-47

158

0.059

1-octanol

C8H17OH

130

-17

194

0.0085

1-nonanol

C9H19OH

144

-6

214

0.00074

1-decanol

C10H21OH

158

+6

233

0.00024

1-dodecanol

C12H25OH

186

+24

259

0.000019

a Deviations from the general trend in melting points occur because melting points for the smallest alcohols are more strongly influenced by differences in crystal structure and lattice energy of the solid. b The properties of 1,5-pentanediol deviate from the trends of the other alcohols because it is a diol and has two -OH groups available for hydrogen bonding. See text.

2. For alcohols of comparable molecular weight, the more hydrogen bonds a compound can form, the more water-soluble the compound, and the higher the boiling and melting points ofthepure compound. In Table 2.3, notice the effect of adding another -OH group to the molecule. The double alcohol 1,5-pentanediol is more water-soluble and has a higher boiling point than single alcohols of comparable molecular weight, as a result of its two -OH groups capable of hydrogen bonding. This effect is general. Double alcohols (diols) are more water-soluble and have higher boiling and melting points than single alcohols of comparable molecular weight. Triple alcohols (triols) are still more water-soluble and have higher boiling and melting points.

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