Hi all,
I've realized I'm still quite a bit confused about enthalpy and am not sure how to think about these problems when I get them.
Does anyone mind confirming that the following is correct?
The standard enthalpy of formation is the energy required to form one mole of compound. The more negative deltaHformation is for a compound, the more energy was released in forming it from its reactant species. Similarly, the more positive deltaHformation is for a compound, the more energy needed to be put into the system to form it--so the bonds broken in the reactants were more stable than the bonds formed to create the compound.
This same idea can be applied to enthalpy of combustion and really any deltaHreaction, right? Since by Hess' Law, deltaHreaction = deltaHformation of products - deltaH formation of reactants. So if deltaHreaction is positive (endothermic rxn), more energy needs to be absorbed by the system to break the bonds in the reactants than energy is released by the system in creating the bonds in the products (so the bonds in the reactants are more stable than the bonds formed in the products). And vice versa for exothermic rxns.
Then with average bond energies, the higher the energy for a bond, the more stable it is (bond dissociation energies are always positive because it takes energy to break a bond). We can determine the enthalpy of formation for a compound also by determining the total energy to break all bonds in the reactants - the total energy to form all bonds in the products to form one mole of the product.
Sorry for the really long post and convoluted thinking. Thank you in advance! Much appreciated.
I've realized I'm still quite a bit confused about enthalpy and am not sure how to think about these problems when I get them.
Does anyone mind confirming that the following is correct?
The standard enthalpy of formation is the energy required to form one mole of compound. The more negative deltaHformation is for a compound, the more energy was released in forming it from its reactant species. Similarly, the more positive deltaHformation is for a compound, the more energy needed to be put into the system to form it--so the bonds broken in the reactants were more stable than the bonds formed to create the compound.
This same idea can be applied to enthalpy of combustion and really any deltaHreaction, right? Since by Hess' Law, deltaHreaction = deltaHformation of products - deltaH formation of reactants. So if deltaHreaction is positive (endothermic rxn), more energy needs to be absorbed by the system to break the bonds in the reactants than energy is released by the system in creating the bonds in the products (so the bonds in the reactants are more stable than the bonds formed in the products). And vice versa for exothermic rxns.
Then with average bond energies, the higher the energy for a bond, the more stable it is (bond dissociation energies are always positive because it takes energy to break a bond). We can determine the enthalpy of formation for a compound also by determining the total energy to break all bonds in the reactants - the total energy to form all bonds in the products to form one mole of the product.
Sorry for the really long post and convoluted thinking. Thank you in advance! Much appreciated.
