Answer:
Lowering the temperature typically reduces the significance of the decrease in entropy. That makes the Gibbs Free energy of the reaction more negative. As a result, the reaction becomes more favorable overall.
Explanation:
In an addition reaction there's a decrease in the number of particles. Consider the hydrogenation of ethene as an example.
[tex]\rm H_2C\text{=}CH_2\; (g) + H_2\; (g) \stackrel{\text{Ni}^\ast}{\to} H_3C\text{-}CH_3\; (g)[/tex].
When [tex]\rm H_2[/tex] is added to [tex]\rm H_2C\text{=}CH_2[/tex] (ethene) under heat and with the presence of a catalyst, [tex]\rm H_3C\text{-}CH3[/tex] (ethane) would be produced.
Note that on the left-hand side of the equation, there are two gaseous molecules. However, on the right-hand side there's only one gaseous molecule. That's a significant decrease in entropy. In other words, [tex]\Delta S < 0[/tex].
The equation for the change in Gibbs Free Energy for a particular reaction is:
[tex]\Delta G = \Delta H + (\underbrace{- T \, \Delta S}_{\text{entropy}\atop \text{term}})[/tex].
For a particular reaction, the more negative [tex]\Delta G[/tex] is, the more spontaneous ("favorable") the reaction would be.
Since typically [tex]\Delta S < 0[/tex] for addition reactions, the "entropy term" of it would be positive. That's not very helpful if the reaction needs to be favorable.
[tex]T[/tex] (absolute temperature) is always nonnegative. However, lowering the temperature could help bring the value of
