Why are there two contradicting formulas for Enthalpy (Delta H): [duplicate]








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[Comparing formula for enthalpy change with bond dissociation energy and formation enthalpy](/questions/10853/comparing-formula-for-enthalpy-change-with-bond-dissociation-energy-and-formatio)

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1) Delta H = Bonds broken - Bonds Formed 


2) Delta H = $H\_{products} - H\_{reactants}$


In the first formula, bonds break in the reactants and form in the products, so its basically saying "bonds energy of reactants" minus the"bond energy of products", which is the complete opposite of what the second formula gives. How so? And when you know to use each formula?




This is my take on this



> 
> 1) Delta H = Bonds broken - Bonds Formed
> 
> 
> 2) Delta H = $H\_{products}−H\_{reactants}$
> 
> 
> 


Equation 2 is pretty straightforward, (change in enthalpy is the enthalpy of the products minus the enthalphy of the reactants).


Equation 1 is a bit less straightforward. Let's define the **system** as where you are having the reaction, and the **surroundings** as everything else. When you break bonds, you take energy **from the surroundings** into the **system**. Then, when bonds are formed, energy is released **from the system** into the **surroundings**. Therefore, if there is more energy released in the bonds formed than there is energy gained in the bonds broken, the **system** loses energy. How I see it is, you are looking from the **surrounding's** perspective at equation 1 and from the **system's** perspective from equation 2, that's why you got two opposite statements (from your POV).


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