Diatomic molecules bond enthalpies

It is relatively simple to assign bond entiialpies to bonds in diatomic molecules. The bond enthalpy is just the energy required to break the diatomic molecule into its component atoms. However, for bonds fliat occur only in polyatomic molecules (such as the C—H bond), we must often utili2e average bond entiialpies. For example, the enthalpy change for tiie following process (called atomization) can be used to define an average bond entiialpy for tiie C—H bond. 

A more useful quantity for comparison with experiment is the heat of formation, which is defined as the enthalpy change when one mole of a compound is formed from its constituent elements in their standard states. The heat of formation can thus be calculated by subtracting the heats of atomisation of the elements and the atomic ionisation energies from the total energy. Unfortunately, ab initio calculations that do not include electron correlation (which we will discuss in Chapter 3) provide uniformly poor estimates of heats of formation w ith errors in bond dissociation energies of 25-40 kcal/mol, even at the Hartree-Fock limit for diatomic molecules. 

TABLE 6.7 Bond Enthalpies of Diatomic Molecules (klmol )... 

All bond enthalpies are positive because heat must be supplied to break a bond. In other words, bond breaking is always endothermic and bond formation is always exothermic. Table 6.7 lists bond enthalpies of some diatomic molecules. 

STRATEGY Decide which bonds are broken and which bonds are formed. Use the mean bond enthalpies in Table 6.8 to estimate the change in enthalpy when the reactant bonds break and the change in enthalpy when the new product bonds form. For diatomic molecules, use the information in Table 6.7 for the specific molecule. Finally, add the enthalpy change required to break the reactant bonds (a positive value) to the enthalpy change that occurs when the product bonds form (a negative value). 

In a diatomic molecule, e.g. chlorine (CI2), the bond dissociation enthalpy and the average bond enthalpy will have the same value. This is because both enthalpy changes refer to the process Cl2(g)--> 2Cl(g). 

Ethyne, ethene and ethane contain 10. 12 and 14 valence electrons, respectively. They are isoelectronic with N2, 02 and F., respectively. Along both series of molecules the central link becomes progressively weaker and the bond lengths increase. The carbon-carbon bond length increases from 121 pm in ethyne to 133 pm in ethene and to 155 pm in ethane. The corresponding bond enthalpy terms are 837, 612 and 348 kJ mol Data for the diatomic molecules is contained in the text. Discuss these data in terms of the VSEPR and MO treatments of the molecules. 

The terms AHAm and AHAy are the enthalpies of atomization of the metal and the nonmetal, respectively. For gaseous diatomic nonmetals, AHA is the enthalpy of dissociation (bond energy plus RT) of the diatomic molecule. For metals which vaporize to form monatomic gases, AHA is identical to the enthalpy of sublimation. If sublimation occurs to a diatomic molecule, M2> then the dissociation enthalpy of the reaction must also he included ... 

Bond enthalpies of diatomic molecules are greater than bond energies by RT. or 3.72 kj-mol-1 at 298 K. 

The strength of a bond, thermochemically speaking, must be measured by the energy required to break it. If we are dealing with a diatomic molecule, this should be given directly by the dissociation enthalpy... 

The relationships between bond enthalpy, bond length and bond order which appear relatively simple in the case of a main group element such as carbon and its compounds, are more difficult to establish when the d-transition metal elements and their compounds are considered. Progress in establishing these relationships for metals is severely hindered by a lack of relevant thermochemical data. This paper reviews some of the more useful information that is available for diatomic molecules, for polynuclear binary carbonyls and for binuclear complexes of the d-transition elements. 

Values of average dissociation energies in text Table 2.4 are actually those for the property of bond enthalpy (see Chapter 6), measured at 298.15 K. The bond dissociation energies of diatomic molecules given in Table 2.3 apply at 0 K (absolute zero). 

Reactants and products in their standard states (pure substance at I bar) at 298.15 K Bond enthalpies for diatomic molecules are given in Table 6.7 in the text. 

In the unique case of diatomic molecules, however, bond energies and dissociation energies are, in principle, identical and exact (although, there may be disagreement in the literature as to what the exact values are). The enthalpy change for the reaction... 

The bond dissociation enthalpy (BDE) for a diatomic molecule AB is the amount of energy necessary to bring about homolytic cleavage, that is, the enthalpy of the dissociation reaction... 

The enthalpy change for the dissociation of a diatomic molecule such as H2 into its atoms in the gas phase can be termed the bond dissociation energy (or more strictly the bond dissociation enthalpy). If we consider the reaction... 

Table 14.3 Some covalent bond enthalpy terms (klmoP ) the values for single bonds refer to the group 15 elements in 3-coordinate environments, and values for triple bonds are for dissociation of the appropriate diatomic molecule.

---