Question

All of the hydrogen halides (H-X) are gaseous in their natural state, but dissolve in water to form the aqueous phase.
Entropies (S) for the gaseous H-X molecules (before reaction) are:
HCl (g): 186.79 J/mol
HI (g): 206.33 J/mol
Entropies (S) for the H-X molecules dissolved/solvated in water (after reaction) are:
HCl (aq): 143.17 J/mol
HI (aq): 219.8 J/mol
1. Calculate the ΔS that each of these H-X compounds undergoes as it transitions from the gas phase to the aqueous phase.

Answer 1

Answer:

The ΔS value of HCl undergoing transitions from gas to the aqueous phase is -43.62 J/mol.

The ΔS value of HI undergoing transitions from gas to the aqueous phase is 13.47 J/mol.

Explanation:

Entropy of the reaction :

$\Delta S=[\text{Sum of entropy of products}]-[\text{Sum of entropy of reactants}]$

1) Entropies the gaseous HCl molecules= $\Delta S_1= 186.79 J/mol$

Entropies the aqueous HCl molecules= $\Delta S_1'= 143.17 J/mol$

$HCl(g)\rightarrow HCl(aq),\Delta S_(g\rightarrow aq)=?$

$\Delta S_(g\rightarrow aq)=\Delta S_1'-\Delta S_1$

$=143.17 J/mol-186.79J/mol=-43.62 J/mol$

The ΔS value of HCl undergoing transitions from gas to the aqueous phase is -43.62 J/mol.

2) Entropies the gaseous HI molecules= $\Delta S_2= 206.33 J/mol$

Entropies the aqueous HI molecules= $\Delta S_2'= 219.8 J/mol$

$HI(g)\rightarrow HI(aq),\Delta S_(g\rightarrow aq)'=?$

$\Delta S_(g\rightarrow aq)'=\Delta S_2'-\Delta S_2$

$=219.8 J/mol-206.33 J/mol=13.47 J/mol$

The ΔS value of HI undergoing transitions from gas to the aqueous phase is 13.47 J/mol.