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ArgumentThere are two common arguments presented as for why (1)$\Delta H < 0$:

Argument (1): Well, indeed I see nothing wrong with the argument presented by the textbook. If adsorption takes place spontaneously, then one can conclude that the change in Gibbs free energy of the process is indeed negative. Since, the entropy change associated with process is necessarily negative  (if we assume the entropy of the adsorbent is necessarily greater in the gaseous or liquid state than it is in the adsorbed state), we need a sufficiently large negative value for the change in enthalpy to ensure spontaneity.

Argument (2)Argument (2): Now, adsorption as a phenomenon is associated with "surface energy" (not, unlike surface tension). The surface of the adsorbent molecule attracts and "attaches" adsorbates either via weak van der waals forces (physisorption) or stronger chemical interactions (chemisorption)--in either cases, the surface energy of the system is minimised due to the formation of these new attractions. Thus, one would say adsorption would be exothermic in nature.

Anyway, this seemsthese seem to be the oft cited reasons. Both these arguments are generally sound and hold true under most, but not all conditions. These arguments are often introduced in most introductory level courses, but they do have some holes.

Food for thought-  -what what if a positive entropy change occurs? Or, the chemical interactions that take place during chemisorption are endothermic in nature? So is adsorption always exothermic, no exceptions? Well, short answer: No.

The following discussion is inspired (largely) by the following publication: http://pubs.acs.org/doi/abs/10.1021/ed038p138?journalCode=jceda8ref 1.

The same need not hold true for chemisorption, since a chemical reaction is indeed taking place. In argument (2), the assumption was that whatever chemical reaction that is taking place at the surface is exothermic, however this need not be the case. If an endothermic reaction takes place then for the process to remain spontaneous the entropy change should not only be positive, but $\Delta ST$$T\Delta S$ must exceed, numerically, $\Delta H$ so that the net Gibbs free energy change as a whole is negative

Summarising, it may be said that in all adsorptions entropy entropy changes due to surface-structural changes in the the adsorbent itself must be considered along with the entropy entropy changes of the adsorbate. In physical adsorption, the the contribution of the first of these two entropy entropy changes to the total entropy change is always likely likely to be insignificant compared to the contribution of of the second (though further investigations alone can establish establish this unambiguously). In chemisorption, however, t,here there is ample evidence already available to show that t,he that the entropy change due to surface structural changes changes in the adsorbent contributes significantly to the the total entropy of adsorption. Whenever the positive entropy entropy change due to adsorbent surface changes exceeds numerically numerically the negative entropy change due to loss loss of freedom of the adsorbate, the net entropy of adsorption adsorption is positive. Failure to recognise this has led led to serious error when using equation (1) to arrive at at the sign of the heat of chemisorption. This, in essence essence, is why endothermic chemisorption has been discounted discounted until very recently, though its existence should should be considered as normal as the existence of endothermic endothermic solution.

You can look up specific examples; the paper I cited references some tooexamples, which are themselves referenced within ref 1. Here's one: http://www.sciencedirect.com/science/article/pii/S0360056408601546

This one One such example is in ref 2, which provides some detailed potential energy diagrams illustrating endothermic chemisorptionschemisorption.


References

(1) Thomas, J. M. The existence of endothermic adsorption. J. Chem. Educ. 1961, 38 (3), 138. DOI: 10.1021/ed038p138.

(2) Dowden, D. A.; Mackenzie, N.; Trapnell, B. M. W. 9 Hydrogen-Deuterium Exchange on the Oxides of Transition Metals. Adv. Catal. 1957, Vol. 9, 65–69. DOI: 10.1016/S0360-0564(08)60154-6.

Argument (1): Well, indeed I see nothing wrong with the argument presented by the textbook. If adsorption takes place spontaneously, then one can conclude that the change in Gibbs free energy of the process is indeed negative. Since, the entropy change associated with process is necessarily negative(if we assume the entropy of the adsorbent is necessarily greater in the gaseous or liquid state than it is in the adsorbed state), we need a sufficiently large negative value for the change in enthalpy to ensure spontaneity.

Argument (2): Now, adsorption as a phenomenon is associated with "surface energy" (not, unlike surface tension). The surface of the adsorbent molecule attracts and "attaches" adsorbates either via weak van der waals forces (physisorption) or stronger chemical interactions (chemisorption)--in either cases, the surface energy of the system is minimised due to the formation of these new attractions. Thus, one would say adsorption would be exothermic in nature.

Anyway, this seems to be the oft cited reasons. Both these arguments are generally sound and hold true under most, but not all conditions. These arguments are often introduced in most introductory level courses, but they do have some holes.

Food for thought--what if a positive entropy change occurs? Or, the chemical interactions that take place during chemisorption are endothermic in nature? So is adsorption always exothermic, no exceptions? Well, short answer: No.

The following discussion is inspired (largely) by the following publication: http://pubs.acs.org/doi/abs/10.1021/ed038p138?journalCode=jceda8

The same need not hold true for chemisorption, since a chemical reaction is indeed taking place. In argument (2), the assumption was that whatever chemical reaction that is taking place at the surface is exothermic, however this need not be the case. If an endothermic reaction takes place then for the process to remain spontaneous the entropy change should not only be positive, but $\Delta ST$ must exceed, numerically, $\Delta H$ so that the net Gibbs free energy change as a whole is negative

Summarising, it may be said that in all adsorptions entropy changes due to surface-structural changes in the adsorbent itself must be considered along with the entropy changes of the adsorbate. In physical adsorption, the contribution of the first of these two entropy changes to the total entropy change is always likely to be insignificant compared to the contribution of the second (though further investigations alone can establish this unambiguously). In chemisorption, however, t,here is ample evidence already available to show that t,he entropy change due to surface structural changes in the adsorbent contributes significantly to the total entropy of adsorption. Whenever the positive entropy change due to adsorbent surface changes exceeds numerically the negative entropy change due to loss of freedom of the adsorbate, the net entropy of adsorption is positive. Failure to recognise this has led to serious error when using equation (1) to arrive at the sign of the heat of chemisorption. This, in essence, is why endothermic chemisorption has been discounted until very recently, though its existence should be considered as normal as the existence of endothermic solution.

You can look up specific examples; the paper I cited references some too. Here's one: http://www.sciencedirect.com/science/article/pii/S0360056408601546

This one provides some detailed potential energy diagrams illustrating endothermic chemisorptions.

There are two common arguments presented as for why $\Delta H < 0$:

Argument (1): Well, indeed I see nothing wrong with the argument presented by the textbook. If adsorption takes place spontaneously, then one can conclude that the change in Gibbs free energy of the process is indeed negative. Since, the entropy change associated with process is necessarily negative  (if we assume the entropy of the adsorbent is necessarily greater in the gaseous or liquid state than it is in the adsorbed state), we need a sufficiently large negative value for the change in enthalpy to ensure spontaneity.

Argument (2): Now, adsorption as a phenomenon is associated with "surface energy" (not, unlike surface tension). The surface of the adsorbent molecule attracts and "attaches" adsorbates either via weak van der waals forces (physisorption) or stronger chemical interactions (chemisorption)--in either cases, the surface energy of the system is minimised due to the formation of these new attractions. Thus, one would say adsorption would be exothermic in nature.

Anyway, these seem to be the oft cited reasons. Both these arguments are generally sound and hold true under most, but not all conditions. These arguments are often introduced in most introductory level courses, but they do have some holes.

Food for thought  - what if a positive entropy change occurs? Or, the chemical interactions that take place during chemisorption are endothermic in nature? So is adsorption always exothermic, no exceptions? Well, short answer: No.

The following discussion is inspired (largely) by ref 1.

The same need not hold true for chemisorption, since a chemical reaction is indeed taking place. In argument (2), the assumption was that whatever chemical reaction that is taking place at the surface is exothermic, however this need not be the case. If an endothermic reaction takes place then for the process to remain spontaneous the entropy change should not only be positive, but $T\Delta S$ must exceed, numerically, $\Delta H$ so that the net Gibbs free energy change as a whole is negative

Summarising, it may be said that in all adsorptions entropy changes due to surface-structural changes in the adsorbent itself must be considered along with the entropy changes of the adsorbate. In physical adsorption, the contribution of the first of these two entropy changes to the total entropy change is always likely to be insignificant compared to the contribution of the second (though further investigations alone can establish this unambiguously). In chemisorption, however, there is ample evidence already available to show that the entropy change due to surface structural changes in the adsorbent contributes significantly to the total entropy of adsorption. Whenever the positive entropy change due to adsorbent surface changes exceeds numerically the negative entropy change due to loss of freedom of the adsorbate, the net entropy of adsorption is positive. Failure to recognise this has led to serious error when using equation (1) to arrive at the sign of the heat of chemisorption. This, in essence, is why endothermic chemisorption has been discounted until very recently, though its existence should be considered as normal as the existence of endothermic solution.

You can look up specific examples, which are themselves referenced within ref 1. One such example is in ref 2, which provides some detailed potential energy diagrams illustrating endothermic chemisorption.


References

(1) Thomas, J. M. The existence of endothermic adsorption. J. Chem. Educ. 1961, 38 (3), 138. DOI: 10.1021/ed038p138.

(2) Dowden, D. A.; Mackenzie, N.; Trapnell, B. M. W. 9 Hydrogen-Deuterium Exchange on the Oxides of Transition Metals. Adv. Catal. 1957, Vol. 9, 65–69. DOI: 10.1016/S0360-0564(08)60154-6.

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However foodFood for thought--what if a positive entropy change occurs? Or, the chemical interactions that take place during chemisorption are endothermic in nature? So is adsorption always exothermic, no exceptions? Well, short answer: No.

The following discussion is inspired largely(largely) by the following publication: http://pubs.acs.org/doi/abs/10.1021/ed038p138?journalCode=jceda8

The same need not hold true for chemisorption, since a chemical reaction is indeed taking place. In argument 2(2), the assumption was that whatever chemical reaction that is taking place at the surface is exothermic, however this need not be the case. If an endothermic reaction takes place then for the process to remain spontaneous the entropy change should not only be positive, but $\Delta ST$ must exceed, numerically, $\Delta H$ so that the net Gibbs free energy change as a whole is negative

However food for thought--what if a positive entropy change occurs? Or, the chemical interactions that take place during chemisorption are endothermic in nature? So is adsorption always exothermic, no exceptions? Well, short answer: No.

The following discussion is inspired largely by the following publication: http://pubs.acs.org/doi/abs/10.1021/ed038p138?journalCode=jceda8

The same need not hold true for chemisorption, since a chemical reaction is indeed taking place. In argument 2, the assumption was that whatever chemical reaction that is taking place at the surface is exothermic, however this need not be the case. If an endothermic reaction takes place then for the process to remain spontaneous the entropy change should not only be positive, but $\Delta ST$ must exceed, numerically, $\Delta H$ so that the net Gibbs free energy change as a whole is negative

Food for thought--what if a positive entropy change occurs? Or, the chemical interactions that take place during chemisorption are endothermic in nature? So is adsorption always exothermic, no exceptions? Well, short answer: No.

The following discussion is inspired (largely) by the following publication: http://pubs.acs.org/doi/abs/10.1021/ed038p138?journalCode=jceda8

The same need not hold true for chemisorption, since a chemical reaction is indeed taking place. In argument (2), the assumption was that whatever chemical reaction that is taking place at the surface is exothermic, however this need not be the case. If an endothermic reaction takes place then for the process to remain spontaneous the entropy change should not only be positive, but $\Delta ST$ must exceed, numerically, $\Delta H$ so that the net Gibbs free energy change as a whole is negative

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You can look up specific examples; the paper I cited references some too. Here's one: http://www.sciencedirect.com/science/article/pii/S0360056408601546

This one provides some detailed potential energy diagrams illustrating endothermic chemisorptions.

You can look up specific examples; the paper I cited references some too.

You can look up specific examples; the paper I cited references some too. Here's one: http://www.sciencedirect.com/science/article/pii/S0360056408601546

This one provides some detailed potential energy diagrams illustrating endothermic chemisorptions.

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