What is true about polar covalent bonds?

Polar Covalent Bonds

A polar covalent bond exists when atoms with different electronegativities share electrons in a covalent bond. Consider the hydrogen chloride (HCl) molecule. Each atom in HCl requires one more electron to form an inert gas electron configuration. Chlorine has a higher electronegativity than hydrogen, but the chlorine atom’s attraction for electrons is not sufficient to remove an electron from hydrogen. Consequently, the bonding electrons in hydrogen chloride are shared unequally in a polar covalent bond. The molecule is represented by the conventional Lewis structure, even though the shared electron pair is associated to a larger extent with chlorine than with hydrogen. The unequal sharing of the bonding pair results in a partial negative charge on the chlorine atom and a partial positive charge on the hydrogen atom. The symbol δ (Greek lowercase delta) denotes these fractional charges.

The hydrogen chloride molecule has a dipole (two poles), which consists of a pair of opposite charges separated from each other. The dipole is shown by an arrow with a cross at one end. The cross is near the end of the molecule that is partially positive, and the arrowhead is near the partially negative end of the molecule.

Single or multiple bonds between carbon atoms are nonpolar. Hydrogen and carbon have similar electronegativity values, so the C—H bond is not normally considered a polar covalent bond. Thus ethane, ethylene, and acetylene have nonpolar covalent bonds, and the compounds are nonpolar.

Bonds between carbon and other elements such as oxygen and nitrogen are polar. The polarity of a bond depends on the electronegativities of the bonded atoms. Large differences between the electronegativities of the bonded atoms increase the polarity of bonds. The direction of the polarity of common bonds found in organic molecules is easily predicted. The common nonmetals are more electronegative than carbon. Therefore, when a carbon atom is bonded to common nonmetal atoms, it has a partial positive charge.

Hydrogen is also less electronegative than the common nonmetals. Therefore, when a hydrogen atom is bonded to common nonmetals, the resulting polar bond has a partial positive charge on the hydrogen atom.

The magnitude of the polarity of a bond is the dipole moment, (D). The dipole moments of several bond types are given in Table 1.2. The dipole moment of a specific bond is relatively constant from compound to compound. When carbon forms multiple bonds to other elements, these bonds are polar. Both the carbon-oxygen double bond in formaldehyde (methanal) and the carbon—nitrogen triple bond in acetonitrile (cyanomethane) are polar.

Table 1.2. Average Dipole Moments (D)

Dipole-Dipole Forces

The bonding electrons in polar covalent bonds are not shared equally, and a bond moment results. However, a molecule may be polar or nonpolar depending on its geometry. For example, tetrachloro-methane (carbon tetrachloride, CCl4) has polar C—Cl bonds, but the tetrahedral arrangement of the four bonds about the central carbon atom causes the individual bond moments to cancel. In contrast, dichloromethane (methylene chloride, CH2Cl2) is a polar molecule with a net polarity away from the partially positive carbon atom toward the partially negative chlorine atoms.

Polar molecules have a negative “end” and a positive “end.” They tend to associate because the positive end of one molecule attracts the negative end of another molecule. The physical properties of polar molecules reflect this association. An increased association between molecules decreases their vapor pressure, which in turn results in a higher boiling point, because more energy is required to vaporize the molecules. The molecular weights and molecular shapes of acetone and isobutane are similar (Figure 2.1), but acetone boils at a higher temperature than isobutane. Acetone contains a polar carbonyl group, whereas isobutane is a nonpolar molecule. The higher boiling point of acetone results from strong the dipole-dipole interaction of the polar carbonyl group.

2.1.1 Unrelaxed permittivity

In the presence of an electric field, electron clouds and atoms in polar covalent bonds may be slightly shifted, inducing a slight polarization that is aligned with the electric field. This acts to store energy and contributes to the capacitive nature of the material. The response times of the electronic and atomic shifts are extremely fast so that, at normal dielectric measurement frequencies, this effect is always present. These induced dipoles are responsible for nonpolar or symmetrically polar polymers having permittivities of 2 or greater. The permittivity due to these induced dipoles is known as the unrelaxed or infinite frequency permittivity (ϵu). At the frequency range typically found in dielectric experiments (10−3 to 108 Hz), induced dipoles react so quickly to an electric field that ϵu is frequency independent.

9.35.5.2.4 Covalent solids

In a good first approximation, and with plausible results, compounds developing polar covalent bonds can often be successfully described by applying ionic, empirical potentials. This is especially true, as long as the local coordination polyhedra of the anions that surround the individual cations are highly symmetric and the overall charge distribution in the solid is approximately isotropic. In such a case, covalent and ionic models in many cases produce the same set of low-energy structure candidates, just with some different energy rankings in the higher-lying minima and possibly different energy barriers between the minima. However, this does not apply to systems where important local minima exhibit different kinds of anisotropic bonding arrangements, for example, sp2- and sp3-hybridized carbon atoms, and one then needs to perform global optimizations on the ab initio energy landscape.

For example, in boron nitride several kinds of, mostly covalent, contributions to the total energy are present, and the global searches need to be performed on the ab initio level. The BN system is particularly interesting as a test system, because the experimentally observed modifications include both layered structures (hexagonal BN) and three-dimensional networks (wurtzite- and sphalerite-type). In global optimizations employing both Hartree–Fock and density functionals,247 all experimentally observed structure types were indeed found. In addition, several new modifications were predicted such as layered structures but with a stacking order different from the experimentally observed structure h-BN. The strength of the general landscape approach has been impressively demonstrated by the discovery of two remarkable new framework structures with low-energies exhibiting the β-BeO structure, and the Al-partial structure in SrAl2, respectively.

Another recent study has employed a combination of data mining, network generation, and local optimization with ab initio energy calculations (DFT), in order to predict crystal structures of group 14 nitrides and phosphides.384 An important step was the generation of new candidates by judiciously substituting atoms of different types in known basic networks, resulting in a multitude of many interesting structures. This procedure is somewhat similar to one of the approaches taken to find crystalline candidates in the Si3B3N7 system.55

12.14.3.5 Fluorescence and Phosphorescence Studies

The fluorescence and absorption spectra of DTT-S,S-dioxide 20a with polar covalent bonds was studied in THF, toluene, and decalin. The spectral line and peak energy are almost independent of the solvent polarity. The fluorescence spectra of the decalin and toluene solutions (almost the same polarity) are red-shifted by about 5 nm, with respect to the THF solution of higher polarity. No evident solvatochromism was observed. The absorbance and fluorescence excitation spectra (at the fluorescence peak wavelength) for DTT-S,S-dioxide 20a (normalized to peak value) was compared. The fluorescence excitation signal is, in fact, dependent both on the density of the excited state (as the absorbance) and on the efficiency of the relaxation from the excited state of the emitting one <2005PCB6004>.

The absorption spectrum of the free DTTPP (3,5-dimethyl-2,6-diphenyl-dithieno[3,2-b;2′,3′-d]thiophene-4,4-dioxide 83 was compared with that of DTTPP bound to antibody anti-CD3. The spectra showed that the absorption of the fluorophore remained unaltered after conjugation, except for some broadening of the low-energy absorption band, probably due to the increase in the number of rotational conformers. Irradiation at λexc = 404 nm led to an intense photoluminescence signal, which retains sizeable intensity when the sample is irradiated at λexc = 480 nm. The significance of this result arises from the fact that 480 nm is the wavelength of the argon-ion laser excitation source of the currently available commercial flow cyclometers. It was found that the activity of the antibody was completely retained in the conjugate, both in the case of the conjugate with antibody anti-CD8 and antibody anti-CD3 <2002CEJ5072>.

Steady-state fluorescence spectra, fluorescence quantum yield (ΦF) and lifetimes (τF) of DTT 15 and DTP 23a were estimated as shown in Table 8. ΦF for DTT is higher than DTP. ΦF for DTP is very small and it was difficult to estimate an accurate fluorescence lifetime by the photon counting method due to weak fluorescence. It is noted that the ΦF for DTP depends largely on the solvent and is 7.7 × 10−5 in acetonitrile. This low ΦF value has been attributed to an addition reaction with the solvent.

Table 8. Peak positions (λF) and quantum yields (ΦF) of fluorescence of bithiophene derivatives

Phosphorescence spectra of DTT and DTP were measured in MeOH/EtOH glass at 77 K. DTT and DTP exhibited phosphorescence with clear vibrational structure. This finding is attributed to the rigid structure caused by the bridging group at the 3,3′ positions of 2,2-bisthiophene 84. The triplet energy of DTT and DTP estimated from the O–O bands of the phosphorescence spectra are given in Table 9.

Table 9. Properties of T–T absorption bands and reaction rates of bithiophene derivatives in triplet excited state

6.1.2.3 Metathesis reactions of AlC and AlH bonds with nonmetallic compounds

The Al

(29)

(30)

(31)

(32)

However, the most common and, in some ways, most useful metathesis of R3Al is with a protic substrate, ROH, R2NH, N

(33)

(34)

(35)

(36)

2.5 Alkyl Halides

This is the first functional class that has a heteroatom. The relatively high electronegativity of the halogens gives a highly polar covalent bond (Inductive effect, Chapter 1). This does not change the sp3 hybrid state or tetrahedral shape of the carbon, but it does give a reactive site that controls the chemistry of alkyl halides. Table 2.6 shows the IUPAC naming of alkyl halides that come from the hydrocarbon parents, with the halogen atom treated as a substituent.

Table 2.6. Selected Common Alkyl Halides

IUPAC, International Union of Pure and Applied Chemistry.

Figure 2.7 shows that, similar to alkanes, alkyl halides and alcohols can be classified as 1° (primary), 2° (secondary), and 3° (tertiary). Note that the nominal oxidation number of the carbon bonded to the halogen changes from −1 in primary to +1 in tertiary. This change explains why there is a difference in reactivity across the range of alkyl halides.

2.1 “Strong” Bonds

Bonds within most organic compounds are described as covalent. In the simplest view of a covalent bond, a pair of electrons is shared by two atoms in the space between them. Each atom formally provides one electron to the bond and these negatively charged electrons are simultaneously attracted to the positive charges of both nuclei. This overcomes the repulsion between the two nuclei. In a polar covalent bond, a part of the electron density of the bonding electron pair is closer to one of the bound nuclei, creating partially positive and negative atomic centers with the magnitude of the charge transfer depending on relative electronegativities of the two atoms.

An ionic bond results after a complete transfer of the bonding electrons from one atom to the other. The resulting positively and negatively charged ions are then electrostatically attracted. Ionic bonds rarely have any particular directionality because they result from electrostatic attraction of each ion to all surrounding ions with opposite charge.

In metallic bonding, bonding electrons are delocalized over a lattice of atoms. In metals, each atom provides one or more electrons that reside between many atomic centers. The free movement of the delocalized electrons results then in important properties of metals such as electrical and thermal conductivity.

13.2.2 Bonding in boron compounds

With the electronic structure of the boron atom being 1s2 2s2 2p1, it might be expected that boron would lose three electrons to give compounds that contain B3+ ions. However, removal of three electrons requires over 6700 kJ mol−1, and this is so high that it precludes compounds that are strictly ionic. Polar covalent bonds exist, and the hybridization can be considered as leading to a set of sp2 hybrid orbitals. However, boron burns readily to produce B2O3, a stable oxide having a heat of formation of −1264 kJ mol−1.

It should be clear that we expect boron to form three equivalent covalent bonds with 120-degree bond angles. As a result, boron halides have the trigonal planar structure (D3h symmetry) for BF3.

In these molecules, the boron atom has only six electrons surrounding it so it interacts readily with species that can function as electron pair donors. For example, when F− reacts with BF3 the product is BF4− in which sp3 hybrids are formed so such species are tetrahedral (Td symmetry). In most cases, molecules containing boron exhibit one of these types of bonding to boron. The boron hydrides represent a special situation that is described later.

9.2 Bonding in Boron Compounds

The electronic structure of the boron atom is 1s2 2s2 2p1. It might be expected that boron would lose the outer electrons and be present in compounds as B3+ ions. This ionization, however, requires over 6700 kJ mol−1, and this amount of energy precludes compounds that are strictly ionic. Polar covalent bonds are much more likely, and it should be kept in mind that orbitals used in bonding are not necessarily the same as the orbitals in isolated atoms. Hybridization of atomic orbitals occurs and in the case of boron, the hybridization can be pictured as follows. Promoting a 2s electron to one of the vacant 2p orbitals can be accomplished with an increase in energy followed by hybridization of the orbitals to produce a set of sp2 hybrid bonding orbitals.

The energy necessary to promote the 2s electron to a 2p level is more than compensated for by the additional energy released when three equivalent bonds are formed.

From the above illustration, we expect boron to form three covalent bonds, which are equal in energy and directed 120° from each other. Accordingly, the boron trihalides, BX3, have the following trigonal planar structure (D3h symmetry).

In fact, all of the compounds containing boron bound to three other atoms have this configuration. In a few cases, such as BH4− and BF4− , sp3 hybrids are formed and the species are tetrahedral (Td symmetry).