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Constants needed: R-8.314 J mol/K, F-96,485 C/mol 1. The rusting of iron is spontaneous and is accompanied by a decrease in the entropy of the entropy change of the system (the iron and oxygen). Why? What can we conclude about the surroundings? (7 pts)
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Answer #1

We know that rusting happens when iron combines with atmospheric oxygen to give hydrated Iron Oxide

2Fe + 3/2 O2 ------------> Fe2O3

In the question it is mentioned that rusting of Iron is a spontaneous process. SO ΔG for the process is >0

Also we see that 1.5 mole of Oxygen is being consumed to give a solid rust mole of iron . Hence the entropy for the reaction is negative ΔS<0 as randomness is decreased.So now we see that entropy of system is <0 , SO entropy of surrounding will increase .SO ΔSsurrounding > 0 (As we know (ΔS +ΔSsurrounding =0)).

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