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Given statement: Enthalpy of adsorption is the strength at which molecule (say $\ce{H2}$) is bound to the adsorbent (say some metal center). For a free $\ce{H2}$ molecule, the entropy contribution is from translational, rotational and vibrational motion. After the adsorption, the translational and rotational motions will not contribute to the entropy change $\mathrm{\Delta S}$. Thus the only contribution to $\mathrm{\Delta S}$ is from the vibrational motion of $\ce{H2}$ which will decrease for increase in adsorption strength.

As per the given statement, increase in enthalpy should decrease the entropy. However, Enthalpy-Entropy correlation states that increase in enthalpy leads to increase in entropy. How this is true?

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closed as unclear what you're asking by Jon Custer, Pritt Balagopal, M.A.R. ಠ_ಠ, NotEvans., Buttonwood Jul 21 '17 at 19:38

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  • $\begingroup$ Could you explain why you think that the enthalpy increases? Isn't the absorption exothermic as, for instance, stated here (--> decrease of enthalpy) ? $\endgroup$ – user45298 Jul 20 '17 at 3:35
  • $\begingroup$ Why, recalling the change of Free / Gibbs enthalpy expressed as $\Delta{}_R G = \Delta_RH - T\Delta_RS$, both changes of enthalpy and of entropy (multiplied by temperature) are taken into consideration? Currently your question reads like one were enough, which I doubt. $\endgroup$ – Buttonwood Jul 21 '17 at 19:38