Bonding in acetylene

SECTION 1.1. VALENCE-BOND APPROACH TO CHEMICAL BONDING 

The hybridization at each carbon atom of acetylene is sp, and the two carbon atoms are considered as bonded by a o--bond and two 7r-bonds, as shown in Fig. 1.4. 

The relation between the number of ligands on carbon (its coordination number), hybridization, and molecular geometry is summarized in Table 1.1. 

Unless all the ligands on a particular carbon atom are identical, there will be deviations from perfectly symmetrical structures. In contrast to methane and carbon tetrachloride, where the bond angles are 109.5°, the C-C-C angle in cyclohexane is 111.5°. The H-C-H angle in formaldehyde is 118°, rather than 120°. Benzene, however, is a regular hexagon with 120° bond angles. 

4 sp Tetrahedral Methane, cyclohexane, methanol, carbon tetrachloride 

---

FIGURE 2 20 Bonding in acetylene based on sp hybridization of carbon The carbon-carbon triple bond is viewed as consisting of one cr bond and two tt bonds... 

As portrayed in Figure 2.20, the two carbons of acetylene are connected to each other by a 2sp-2sp a bond, and each is attached to a hydrogen substituent by a 2sp- s a bond. The unhybridized 2p orbitals on one carbon overlap with their counterparts on the other to fonrr two tt bonds. The carbon-carbon triple bond in acetylene is viewed as a multiple bond of the a + tt + tt type. 

An sp hybridization model for the caibon-caibon triple bond was developed in Section 2.21 and is reviewed for acetylene in Figure 9.2. Figure 9.3 compares the electrostatic potential maps of ethylene and acetylene and shows how the second tt bond in acetylene causes a band of high electron density to encircle the molecule. 

The length of the carbon-carbon triple bond in acetylene is 120 pm, and the strength is approximately 835 kj/mol (200 kcal/mol), making itthe shortest and strongest known carbon-carbon bond. Measurements show that approximately 318 kj/mol (76 kcal/mol) is needed to break a rbond in acetylene, a value some 50 kj/mol larger than the 268 kj/mol needed to break an alkene n bond. 

With regard to the different points of view outlined in (a), (b) and (c), it should be pointed out that these differences arise mainly from the use of localized (a, LMO), or canonical (CMO, b, and c) molecular orbitals. In principle LMOs and CMOs are equivalent and are related by a unitary transformation. This can be illustrated by the C=C bonding in acetylene. 

Ervin et al. [27] have determined the electron affinity of the acetylide radical, HC = C-, to be equal to 2.969 + 0.010 eV and the enthalpy of the acid dissociation of acetylene in the gas phase to be equal to 377.8 + 0.6 Kcal mol Use these data, together with the ionization potential of the hydrogen atom, 13.595 eV, to calculate the enthalpy for the dissociation of the CH bond in acetylene. The ionization potentials are properly applied at 0 K, but a good approximation is to assume that they are equal to enthalpy changes at 298.15 K, the temperature at which the enthalpy of the acid dissociation was measured. 

The bonding in acetylene has one C-C a bond together with two C-H a bonds the p orbitals on each carbon, each carrying one electron, interact by side-to-side overlap to produce two Jt bonds (Figure 2.19). Note again that the p orbitals can only overlap if their axes are parallel. 

Acetylene first became the primary intermediate molecule in chemical synthesis in the early twentieth century. As noted previously, the triple bond in acetylene is ideal to add one or two small molecules. For example, addition of one mole of water produces vinyl alcohol,... 

The experimental data span a very wide range from the CH bond in acetylene (27 kcal/mol stronger than that in methane) to the CH bond in cycloheptatriene (31 kcal/mol weaker). This presumably reflects the stability (in instability) of the radical product more than it does the hydrocarbon reactant. The usual models have been surveyed. 

Problem 8.54 Is there any inconsistency between the facts that the C—H bond in acetylene has the greatest bond energy of all C—H bonds and that it is also the most acidic <... 

The triple bond in acetylene results in a high energy content that is released when acetylene is burned. After Willsons discovery of a method to produce commercial quantities of... 

The carbon-carbon triple bond in acetylene can be treated in a similar way to that in the nitrogen molecule.8 The details of hybridisation may... 

Carbon can form multiple covalent bonds by sharing more than two electrons with a neighboring atom (Section 7.5). In ethylene, the two carbon atoms share four electrons in a double bond. In acetylene, the two carbons share six electrons in a triple bond ... 

The flow of electrons from a Lewis base to a Lewis acid is often indicated with a curved arrow. The arrow starts on a pair of nonbonding electrons on the Lewis base and points toward the Lewis acid with which it reacts. Because adding a pair of electrons to one point on a molecule often displaces electrons in the molecule, combinations of curved arrows are often used to describe even simple chemical reactions. Consider the following example, in which a pair of electrons on an NH2 ion are donated to the H+ ion formed when the electrons in one of the CH bonds in acetylene are given to the carbon atom instead of being shared by the C and H atoms in this bond. 

How atoms are bonded to each other plays a significant role in the properties of the substances they create. Saturated and unsaturated fats are listed on the nutritional labels of many food products, and is medical advice about the advisability of including them in a healthy diet is offered. Saturated fats have single bonds, whereas unsaturated fats have double bonds. The triple bond in acetylene is responsible for its extreme reactivity and flammability. 

Electron-poor multiple bonds in acetylene dicarboxylates42 137 149), ketenes 150 152> isocyanates, isothiocyanates, carbodiimides 138-153) and arynes42) add metallated ynamines to give adducts that still carry the ynamine moiety. This behaviour contrast with that of other ynamines which mostly undergo various cycloadditions but it must be stressed that even here cycloadducts may be formed as by-products. 

The triple bond is relatively short because of the attractive overlap of three bonding pairs of electrons and the high s character of the sp hybrid orbitals. The sp hybrid orbitals are about one-half s character (as opposed to one-third s character of sp2 hybrids and one-fourth of sp3 hybrids), using more of the closer, tightly held s orbital. The sp hybrid orbitals also account for the slightly shorter C — H bonds in acetylene compared with ethylene. 

For example, the central bond in acetylene is a triple bond H-C-C-H. 

That the C—H bond in acetylene (and HCN) deviates from that in methane is seen from the short length of 1.059 A (or 1.057 A) compared with 1.093 A, and from the positive charge of the hydrogen atom in both cases, see p. 225 (ethylene C—H... 

The discrepant behaviour of the C—H bond in acetylene shows itself in the differences which are apparently found for the C=C bond energy (when the C—H and C—C bond energies are taken to be the same throughout ) for acetylene, methyl-and dimethyl-acetylene ofi 28.6kcal, 132.2 kcal and 135.5 kcal. 

The triple bond in acetylene is seen to consist of one c bond joining the line-of-centers between the two carbon atoms, and two n bonds whose lobes of electron density are in mutually-perpendicular planes. The acetylene molecule is of course linear, since the angle between the two sp hybrid orbitals that produce the c skeleton of the molecule is... 

---