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    ΔHf of C6H5CH2CH3 = -48.95 Kj/ mol

    ΔHf of SO2 = -296 Kj/ mol

    ΔHf of SiO2 = − 910.86 Kj/ mol

    Is this a standard state of Enthalpy correct? I'm confused because i'm not sure whether i need to add an oxygen, balance and then use hess in order to get ΔH combustion for the above elements.

    I am supposed to just get ΔH combustion figures for the above, i found the above figures in a Standard enthalpy change of formation (data table) - Wikipedia, the free encyclopedia.

    I know the (-) identifies an exothermic reaction.

    Any advice would be appreciated.


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    (Original post by g-wizz)
    ΔHf of C6H5CH2CH3 = -48.95 Kj/ mol

    ΔHf of SO2 = -296 Kj/ mol

    ΔHf of SiO2 = − 910.86 Kj/ mol

    Is this a standard state of Enthalpy correct? I'm confused because i'm not sure whether i need to add an oxygen, balance and then use hess in order to get ΔH combustion for the above elements.

    I am supposed to just get ΔH combustion figures for the above, i found the above figures in a Standard enthalpy change of formation (data table) - Wikipedia, the free encyclopedia.

    I know the (-) identifies an exothermic reaction.

    Any advice would be appreciated.


    The enthalpy of combustion of the bottom two is the same as the enthalpy of formation.
    If you wrote out a formation equation and a combustion equation this is immediately apparent . . .
    I don't actually know if silicon combusts at all
 
 
 
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