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Chemistry Forums for Students => Undergraduate General Chemistry Forum => Topic started by: aaron-ch on August 08, 2012, 04:28:53 PM

Title: Oxidation of Iron
Post by: aaron-ch on August 08, 2012, 04:28:53 PM

Hi We were given the problem:

"Write the chemical equation for the oxidation of Fe and of Al.  Use ΔG_f^° to determine whether either process is spontaneous at 25°C"

I am really confused with how I would write the oxidation equations with out being given information about how Fe and Al ionize, and what they react with?

please help...   

Title: Re: Oxidation of Iron
Post by: discodermolide on August 09, 2012, 01:24:56 AM

Quote from: aaron-ch on August 08, 2012, 04:28:53 PM

Hi We were given the problem:

"Write the chemical equation for the oxidation of Fe and of Al.  Use ΔG_f^° to determine whether either process is spontaneous at 25°C"

I am really confused with how I would write the oxidation equations with out being given information about how Fe and Al ionize, and what they react with?

please help...   

The key word here is oxidation!
So Fe and Al are reacting with oxygen.
So now you can write the balanced equations.

Title: Re: Oxidation of Iron
Post by: mugabo daniel on August 09, 2012, 04:06:48 PM

iron undergoes oxidation to form two compounds i.e. iron(ii) oxide and iron(iii) oxide depending on oxygen content.
    2Fe + O2 --> 2FeO
    4Fe + 3O2 --> 2Fe2O3
Try to find energy changes for the above reactions and substitute in the Gibbs formula;
      DG = DH-TDS
  where H = enthalpy change
           T = temperature in kelvin
           S = entropy change
Aluminium undergoes oxidation to form aluminium oxide
   2Al + 3O2--> Al2O3
if G is negative then the reaction is spontaneous and viceversa.
this is what i know u can correct me if am wrong.