What is the heat of hydrogenation in a hydrogenation reaction?

Hydrogenation reactions consist of the addition of (guess what?) hydrogen to a molecule. For example...

#"Ethene + "H_2 "" stackrel("Pd/C")(->) "Ethane"#

The heat of whatever event at constant pressure, #q_p#, is simply the enthalpy of such an event, #DeltaH#. For a hydrogenation reaction, the enthalpy of hydrogenation is simply the enthalpy of reaction, or #DeltaH_"rxn"#.

This enthalpy could be broken down into which bonds were broken or made. One might call those #DeltaH_"broken"# and #DeltaH_"made"#.

Whatever the case, the heat of hydrogenation is fundamentally based on which bonds were broken, which were made, and the overall differences in them throughout a hydrogenation reaction typically on a per-#"mol"# basis, commonly of alkenes, sometimes of alkynes.

In the example I listed above, you break:

• #1# #C=C# bond

and make:

• #1# #C-C# bond

since you had a double bond and then you just have a single bond. Then you break:

• #1# #H-H# bond

before you allow #H_2# to make:

• #2# #C-H# bonds

These enthalpies are:

#C=C# bond: #"~602 kJ/mol"#
#C-C# bond: #"~346 kJ/mol"#
#H-H# bond: #"~436 kJ/mol"#
#C-H# bond: #"~413 kJ/mol"#

Breaking a bond takes outside energy and puts it into the bond, and is thus positive. Making a bond releases energy into the atmosphere and is thus reported as negative. Overall, you get about:

#DeltaH_"rxn" = sum_i DeltaH_"broken,i" - sum_i DeltaH_"made,i"#

#= overbrace([(underbrace(602)_(broken) - underbrace(346)_(made)) + underbrace(436)_(broken)])^("Step 1") - overbrace([underbrace(413 + 413)_(made)])^("Step 2")#

#= overbrace((602+436))^(broken) - overbrace((346 + 413 + 413))^(made)#

# ~~ color(blue)("-134 kJ/mol")#

The enthalpy or heat of hydrogenation of ethene into ethane is exothermic.