Formation reactions enthalpy

Sfi f-Test 6.13B Methanol is a clean-burning liquid fuel proposed as a replacement for gasoline. Suppose it could be produced by the controlled reaction of the oxygen in air with methane. Find the standard reaction enthalpy for the formation of I mol CHjOH(l) from methane and oxygen, given the following information ... 

There are millions of possible reactions, and it is impractical to list every one with its standard reaction enthalpy. However, chemists have devised an ingenious alternative. First, they report the standard enthalpies of formation of substances. Then they combine these quantities to obtain the standard enthalpy of reaction needed. Let s look at these two stages in turn. 

The standard enthalpy of formation, AH°, of a substance is the standard reaction enthalpy per mole of formula units for the formation of a substance from its elements in their most stable form, as in the reaction... 

Now let s see how to combine standard enthalpies of formation to calculate a standard reaction enthalpy. To do so, we imagine carrying out the reaction in two steps we reverse the formation of the reactants from the elements, then combine the elements to form the products. The first step is usually to calculate the reaction enthalpy for the formation of all the products from their elements. For this step, we use the enthalpies of formation of the products. Then, we calculate the reaction enthalpy for the formation of all the reactants from their elements. The difference between these two totals is the standard enthalpy of the reaction (Fig. 6.31) ... 

FIGURE 6.31 The reaction enthalpy can be constructed from enthalpies of formation by imagining forming the reactants and products from their elements. Then the reaction enthalpy is the difference between the enthalpies of products and reactants. 

STRATEGY We expect a strongly negative value because all combustions are exothermic and this oxidation is like an incomplete combustion. First, add up the individual standard enthalpies of formation of the products, multiplying each value by the appropriate number of moles from the balanced equation. Remember that the standard enthalpy of formation of an element in its most stable form is zero. Then, calculate the total standard enthalpy of formation of the reactants in the same way and use Eq. 20 to calculate the standard reaction enthalpy. 

Standard enthalpies of formation are commonly determined from combustion data by using Eq. 20. The procedure is the same, but the standard reaction enthalpy is known and the unknown value is one of the standard enthalpies of formation. 

A note on good practice. The use of mean bond enthalpies is hazardous because actual bond enthalpies often differ considerably from mean values. The modem procedure for estimating a reaction enthalpy is to use commercial software to calculate the enthalpies of formation of the reactants and products and then to take the difference, as in Section 6.18. 

Calculate the reaction enthalpy for the formation of anhydrous aluminum chloride, 2 Al(s) + 3 CI2(g) —> 2 AICI3(s), from the following data ... 

Using standard enthalpies of formation from Appendix 2A, calculate the standard reaction enthalpy for each of the following reactions ... 

Just as we can combine standard enthalpies of formation to obtain standard reaction enthalpies, we can also combine standard Gibbs free energies of formation to obtain standard Gibbs free energies of reaction ... 

STRATEGY Raising the temperature of an equilibrium mixture will tend to shift its composition in the endothermic direction of the reaction. A positive reaction enthalpy indicates that the reaction is endothermic in the forward direction. A negative reaction enthalpy indicates that the reaction is endothermic in the reverse direction. To find the standard reaction enthalpy, use the standard enthalpies of formation given in Appendix 2A. 

Standard reaction enthalpy (X = H) and Gibbs free energy (X = G) from standard enthalpies of formation ... 

The addition of an ion to butadiene is clearly an exothermic process in the gas phase due to the formation of aa-bond substituting a rc-bond. The agreement of the reaction enthalpies of the reactions (11) and (12) with equal R (except R = H) is surprising (Table 11). 

All alkyl ions tested demonstrate a comparable behaviour independent of the sign of their charges. The decrease of the reaction enthalpies AH (11) with the change from the methyl to the ethyl cation (AAH (ll) = 165 kJ mol-1) and from the ethyl to the but-2-enyl cation (AAH°(11) = 117 kJ mol-1) corresponds to the increase of stability of these carbenium ions, which are expressed by the difference of their heats of formation (AAH f = —118 and AAHj = —42 kJ mol-1 90)) and of their hydride ion affinity (AHIA = 176 and 126 kJ mol-1 91)), respectively. 

Using standard enthalpies from Ref. 132) a reaction enthalpy of 92 kJ mol-1 is the result for reaction (24). This value lies near the above estimated limit and not far from the values which have been calculated from experimental heats of formation for the second and third propagation step (102 kJ mol-1). 

Thermal dimerization of ethylene to cyclobutane is forbidden by orbital symmetry (Sect 3.5 in Chapter Elements of a Chemical Orbital Theory by Inagaki in this volume). The activation barrier is high E =44 kcal mof ) [9]. Cyclobutane cannot be prepared on a preparative scale by the dimerization of ethylenes despite a favorable reaction enthalpy (AH = -19 kcal mol" ). Thermal reactions between alkenes usually proceed via diradical intermediates [10-12]. The process of the diradical formation is the most favored by the HOMO-LUMO interaction (Scheme 25b in chapter Elements of a Chemical Orbital Theory ). The intervention of the diradical intermediates impfies loss of stereochemical integrity. This is a characteric feature of the thermal reactions between alkenes in the delocalization band of the mechanistic spectrum. 

How do we determine the energy and enthalpy changes for a chemical reaction We could perform calorimetry experiments and analyze the results, but to do this for every chemical reaction would be an insurmountable task. Furthermore, it turns out to be unnecessary. Using the first law of thermodynamics and the idea of a state function, we can calculate enthalpy changes for almost any reaction using experimental values for one set of reactions, the formation reactions. 

When the partial pressure of each gaseous reagent is 1 bar and the concentration of each species in solution is 1 M, the conditions are defined to be standard. Under these conditions, the enthalpy change in a formation reaction is the standard enthalpy of formation (A... 

The enthalpy change of the formation reaction is just A of N2 O4 A fojmation = A //f (N2 O4) Thus, the molar enthalpy change of the overall reaction can be expressed entirely in terms of standard formation reactions A - r°eaction = A (N2 O4) - 2 A i7f (N O2) The enthalpy change of the overall reaction is the sum of the formation enthalpies of the products minus the sum of the formation enthalpies of the reactants. 

Our analysis of the reaction of nitrogen dioxide molecules is not unique. The same type of path can be visualized for any chemical reaction, as Figure 6-20 shows. The reaction enthalpy for any chemical reaction can be found from the standard enthalpies of formation for all the reactants and products. Multiply each standard enthalpy of formation by the appropriate stoichiometric coefficient, add the values for the products, add the values for the reactants, and subtract the sum for reactants from the sum for products. Equation summarizes this procedure ... 

Equation can also be used to calculate the standard enthalpy of formation of a substance whose formation reaction does not proceed cleanly and rapidly. The enthalpy change for some other chemical reaction involving the substance can be determined by calorimetric measurements. Then Equation can be used to calculate the unknown standard enthalpy of formation. Example shows how to do this using experimental data from a constant-volume calorimetry experiment combined with standard heats of formation. 

Again, it is convenient to follow the seven-step procedure to solve this problem. We are asked to find an enthalpy of formation. Because enthalpy is a state function, we can visualize the reaction as occurring through decomposition and formation reactions. Appendix D lists enthalpies of formation, and the experimental heat of combustion is provided. We can use Equation to relate the enthalpy of combustion to the standard enthalpy of formation for octane. 

Entropy changes are important in every process, but chemists are particularly interested in the effects of entropy on chemical reactions. If a reaction occurs under standard conditions, its entropy change can be calculated from absolute entropies using the same reasoning used to calculate reaction enthalpies from standard enthalpies of formation. The products of the reaction have molar entropies, and so do the reactants. The total entropy of the products is the sum of the molar entropies of the products multiplied by their stoichiometric coefficients in the balanced chemical equation. The total entropy of the reactants is a similar sum for the reactants. Equation... 

Calorimetry investigations of zinc ions with functionalized pyridines have been carried out in both dimethylformamide and acetonitrile. The pyridines used were pyridine, 3-methylpyridine, and 4-methylpyridine. In DMF, for all three pyridines, four- and six-coordinate species formed and their formation constants, reaction enthalpies and entropies were determined. The stability increases linearly with increasing basicity of the pyridine derivative. The formation of the 3-methylpyridine complex is enthalpically less favorable and entropically more favorable than... 

We can calculate the value of A// tn for the cast-formation reaction at 298 K using the enthalpies of formation of reactants and products ... 

KEY TERMS endothermic reaction enthalpy of reaction dissolution enthalpy of formation... 

The standard deviation between experimental and calculated heats of reaction are between 0.5 and 1 kcal/mol for those classes of compounds where enough experimental heats of formation are available to allow a full parameterization. For those classes of compounds where insufficient heats of formation are known to allow the determination of all parameters for 1,2- and 1,3-interactions, an estimate can be given for the bond energy terms which are the dominating parameters. Even here, therefore, a reasonable value for the reaction enthalpy is available. 

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