What is the enthalpy of hydrogenation?

Well, it is just like any other reaction... you could simply find the enthalpy of reaction for a particular hydrogenation reaction (i.e. addition of #"H"_2(g)# across a #pi# bond), such as...

The standard enthalpy of hydrogenation of ethene would be gotten from...

#DeltaH_(rxn)^@ = sum_P nu_P DeltaH_(f,P)^@ - sum_R nu_R DeltaH_(f,R)^@#

where:

• #nu# is the stoichiometric coefficient.
• #R# and #P# stand for reactant and product, respectively.
• #DeltaH^@# is the standard molar enthalpy.

With enthalpy of formation data obtained from NIST, we have:

#DeltaH_(f,"C"_2"H"_4(g))^@ = "52.47 kJ/mol"#

#DeltaH_(f,"H"_2(g))^@ = "0.00 kJ/mol"# (why?)

#DeltaH_(f,"C"_2"H"_6(g))^@ = -"83.8 kJ/mol"#

We get...

#color(blue)(DeltaH_(hydr)^@)#

#= [nu_(C_2H_6(g))DeltaH_(f,"C"_2"H"_6(g))^@] - [nu_(C_2H_4(g))DeltaH_(f,"C"_2"H"_4(g))^@ + nu_(H_2(g))DeltaH_(f,"H"_2(g))^@]#

#= ["1 equiv." xx overbrace(-"83.8 kJ/mol")^(DeltaH_(f,"C"_2"H"_6(g))^@)] - ["1 equiv." xx overbrace("52.47 kJ/mol")^(DeltaH_(f,"C"_2"H"_4(g))^@) + "1 equiv." xx overbrace("0.00 kJ/mol")^(DeltaH_(f,"H"_2(g))^@)]#

#= color(blue)(-"136.27 kJ/mol")#