Chemical reactions can be used to produce flames for heating. The nonnuclear reaction with the highest attainable flame temperature (approximately 6000°C, about the surface temperature of the sun) is that between hydrogen and fluorine: [latex]H_{2}(g) + F_{2}(g) → 2\ HF(g)[/latex] Use the molar bond enthalpies in Table 14.5 to estimate the value of [latex]\Delta H^{\circ}_{rxn}[/latex] for this equation. Compare this to the value calculated using the molar enthalpies of formation in Table 14.3.

The reaction involves the breaking of one hydrogen–hydrogen bond and one fluorine–fluorine bond and the formation of two hydrogen– fluorine bonds. Thus, [latex]ΔH^{\circ}_{rxn} ≈ H_{bond}(H–H) + H_{bond} (F–F) – 2\ H_{bond} (H–F)\\= (435\ kJ·mol^{–1}) + (155\ kJ·mol^{–1}) – (2)(565\ kJ·mol^{–1}) \\ = – 540\ kJ·mol^{–1}[/latex] Because this reaction is the formation reaction for two moles of [latex]HF(g)[/latex] from its constituent elements, the experimentally determined value of [latex]\Delta H^{\circ}_{rxn}[/latex] is twice the molar enthalpy of formation of [latex]HF(g)[/latex] listed in Table 14.3, or [latex]– 546\ kJ·mol^{–1}[/latex] , a value with less than a 2% difference from that which we found using average molar bond enthalpies.

Question 14.9: Chemical reactions can be used to produce flames for heating......

General Chemistry [1096620]

Chemical reactions can be used to produce flames for heating. The nonnuclear reaction with the highest attainable flame temperature (approximately 6000°C, about the surface temperature of the sun) is that between hydrogen and fluorine:

H_{2}(g) + F_{2}(g) → 2\ HF(g)

Use the molar bond enthalpies in Table 14.5 to estimate the value of $\Delta H^{\circ}_{rxn}$ for this equation. Compare this to the value calculated using the molar enthalpies of formation in Table 14.3.

TABLE 14.3 standard enthalpies of formation, $ΔH^{\circ}_{f}$ , for various substances a 25°c
Substance	Formula	$ΔH^{\circ}_{f}/kJ\cdot mol^{-1}$	Substance	Formula	$ΔH^{\circ}_{f}/kJ\cdot mol^{-1}$
aluminum oxide	$Al_{2}O_{3}(s)$	-1675.7	hydrogen fluoride	$HF(g)$	-273.3
ammonia	$NH_{3}(g)$	-45.9	hydrogen iodide	$HI(g)$	+26.5
benzene	$C_{6}H_{6}(l)$	+49.1	hydrogen peroxide	$H_{2}O_{2}(l)$	-187.8
benzoic acid	$C_{6}H_{5}COOH(s)$	-385.2	iodine vapor	$I_{2}(g)$	+62.4
bromine vapor	$Br_{2}(g)$	+30.9	magnesium carbonate	$MgCO_{3}(s)$	-1095.8
butane	$C_{4}H_{10}(g)$	-125.7	magnesium oxide	$MgO(s)$	-601.6
calcium carbonate	$CaCO_{3}(s)$	-1207.6	magnesium sulfide	$MgS(s)$	-346.0
carbon (diamond)	$C(s)$	+1.897	methane	$CH_{4}(g)$	-74.6
carbon (graphite)	$C(s)$	0	methanol (methyl alcohol)	$CH_{3}OH(l)$ $CH_{3}OH(g)$	-239.2 -201.0
carbon (buckminster fullerene)	$C_{60}(s)$	+2327.0	methyl chloride	$CH_{3}Cl(g)$	-81.9
carbon dioxide	$CO_{2}(g)$	-393.5	nitrogen dioxide	$NO_{2}(g)$	+33.2
carbon monoxide	$CO(g)$	-110.5	nitrogen oxide	$NO(g)$	+91.3
carbon tetrachloride	$CCl_{4}(l)$ $CCl_{4}(g)$	-128.2 -95.7	dinitrogen tetroxide	$N_{2}O_{4}(g)$ $N_{2}O_{4}(l)$	+11.1 -19.5
chromium (III) oxide	$Cr_{2}O_{3}(s)$	-1139.7	octane	$C_{8}H_{18}(l)$	-250.1
cyclohexane	$C_{6}H_{12}(l)$	-156.4	pentane	$C_{5}H_{12}(l)$	-173.5
ethane	$C_{2}H_{6}(g)$	-84.0	propane	$C_{3}H_{8}(g)$	-103.8
ethanol (ethyl alcohol)	$CH_{3}CH_{2}OH(l)$	-277.6	sodium carbonate	$Na_{2}CO_{3}(s)$	-1130.7
ethene (ethylene)	$C_{2}H_{4}(g)$	+52.4	sodium oxide	$Na_{2}O(s)$	-414.2
ethyne (acetylene)	$C_{2}H_{2}(g)$	+227.4	sucrose	$C_{12}H_{22}O_{11}(s)$	-2226.1
freon-12 (dichloro difluoromethane)	$CF_{2}Cl_{2}(g)$	-477.4	sulfur dioxide	$SO_{2}(g)$	-296.8
glucose	$C_{6}H_{12}O_{6}(s)$	-1273.3	sulfur trioxide	$SO_{3}(g)$	-395.7
hexane	$C_{6}H_{14}(l)$	-198.7	tin(IV) oxide	$SnO_{2}(s)$	-577.6
hydrazine	$N_{2}H_{4}(l)$ $N_{2}H_{4}(g)$	+50.6 +95.4	water	$H_{2}O(l)$ $H_{2}O(g)$	-285.8 -241.8
hydrogen bromide	$HBr(g)$	-36.3
hydrogen chloride	$HCl(g)$	-92.3
Data from CRC Handbook of Chemistry and Physics, 86th Ed., Ed. David R. Lide, CRC Press, 2005–2006. (More thermodynamic data are given in Appendix D.)

TABLE 14.5 average molar bond enthalpies
Bond	Molar bond enthalpy, $H_{bond}/kJ\cdot mol^{-1}$	Bond	Molar bond enthalpy, $H_{bond}/kJ\cdot mol^{-1}$
O–H	464	C≡N	890
O–O	142	N–H	390
C–O	351	N–N	159
O=O	502	N=N	418
C=O	730	N≡N	945
C–C	347	F–F	155
C=C	615	Cl–Cl	243
C≡C	811	Br–Br	192
C–H	414	H–H	435
C–F	439	H–F	565
C–Cl	331	H–Cl	431
C–Br	276	H–Br	368
C–N	293	H–S	364
C=N	615

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The reaction involves the breaking of one hydrogen–hydrogen bond and one fluorine–fluorine bond and the formation of two hydrogen– fluorine bonds. Thus,

ΔH^{\circ}_{rxn} ≈ H_{bond}(H–H) + H_{bond} (F–F) – 2\ H_{bond} (H–F)\\= (435\ kJ·mol^{–1}) + (155\ kJ·mol^{–1}) – (2)(565\ kJ·mol^{–1}) \\ = – 540\ kJ·mol^{–1}

Because this reaction is the formation reaction for two moles of $HF(g)$ from its constituent elements, the experimentally determined value of $\Delta H^{\circ}_{rxn}$ is twice the molar enthalpy of formation of $HF(g)$ listed in Table 14.3, or $– 546\ kJ·mol^{–1}$ , a value with less than a 2% difference from that which we found using average molar bond enthalpies.