BR Organic Chemistry Chapter 1 Enthalpy

[dorian baltar]

In BR O Chem Chapter 1, enthalpy is discussed and the equation [delta H = The sum (Energy of bonds broken - Energy of bonds formed)] is given.
The equation seems incorrect to me and I think formed and broken should be switched.
If the bonds formed (the bonds in the product) have more energy than those broken in the reactants, then there would have to be a net input of energy into the reaction, and the reaction would be endothermic. However, numbers like that would give a negative number for enthalpy, which is exothermic.
It might be the way it's stated, because I get Hf-Hi = delta H. I take that to mean that the sum of the energy of all the bonds at the end of the reaction minus the sum of the energy of all the bonds before the reaction is the change in enthalpy.
Am I just looking at this in a wrong way?
It's been bugging me for over a week.

I greatly appreciate any help!!!

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[KeepingitMoving]

Try looking at it this way...enthalpy is simply the difference between the sum of the enthalpies of the products and the sum of the enthalpies of the reactants.
Remember that breaking bonds and forming bonds BOTH require energy.

• If the reactants have weak bonds, while the products have strong bonds then the reaction is exothermic (enthalpy change < 0). There is a small amount of energy needed to break the bond (smaller bond energy) and a bigger energy released when strong bonds form. A negative enthalpy change means that the system released energy.

• If the reactants have strong bonds, but the products have weak bonds its an endothermic reaction. (enthalpy change > 0). The energy required to break the reactant bonds is greater than the energy released when the product bonds form.

I got this information online...check out the website for more detail
http://chemwiki.ucdavis.edu/Theoretical_Chemistry/Chemical_Bonding/General_Principles/Bond_Energies

 

[Jepstein30]

In BR O Chem Chapter 1, enthalpy is discussed and the equation [delta H = The sum (Energy of bonds broken - Energy of bonds formed)] is given.
The equation seems incorrect to me and I think formed and broken should be switched.
If the bonds formed (the bonds in the product) have more energy than those broken in the reactants, then there would have to be a net input of energy into the reaction, and the reaction would be endothermic. However, numbers like that would give a negative number for enthalpy, which is exothermic.
It might be the way it's stated, because I get Hf-Hi = delta H. I take that to mean that the sum of the energy of all the bonds at the end of the reaction minus the sum of the energy of all the bonds before the reaction is the change in enthalpy.
Am I just looking at this in a wrong way?
It's been bugging me for over a week.

I greatly appreciate any help!!!

sounds like to me that you're comparing bond enthalpies rather than considering the enthalpy change of a reaction. There's a distinction there.. you talked about the product's bonds having "more energy" than the broken bonds.. but remember, them having MORE energy means they would RELEASE more energy when they are formed! You're using the wrong sign for the "energy of bonds formed" portion of the equation.

in general,
breaking bonds = requires energy
forming bonds = releases energy

in a reaction, we need to break the reactant's bonds.. and form the product's bonds..
so we have 2 scenarios:
1) it takes more energy to break the reactant's bonds than energy that is released when forming the product's bonds -> we need more energy to do the reaction so this is endothermic
2) it takes less energy to break the reactant's bonds than energy that is released when forming the product's bonds -> we have "leftover" energy so this is exothermic

in these scenarios, we can simply compare the sum of the relevant bond enthalpies on each side.
[delta H = The sum (Energy of bonds broken - Energy of bonds formed)]

same scenarios again:
1) the "energy of bonds broken" is larger than the "energy of bonds formed" so we get a positive number = endothermic
2) the "energy of bonds broken" is smaller than the "energy of bonds formed" so we get a negative number = exothermic

if we switch it around, we'd get the opposite result.

 

[dorian baltar]

So can I qualitatively think about change in enthalpy by considering the distances between the energy of the reactant and product and the activation energy? For example, in the combustion of a hydrocarbon, the energy required to reach the activation energy is less than the energy required for the product to reach the activation energy.
If that's true, great. It's really easy for me to visualize a reaction coordinate.

 

[Jepstein30]

So can I qualitatively think about change in enthalpy by considering the distances between the energy of the reactant and product and the activation energy? For example, in the combustion of a hydrocarbon, the energy required to reach the activation energy is less than the energy required for the product to reach the activation energy.
If that's true, great. It's really easy for me to visualize a reaction coordinate.

I think that works. Never heard of it being compared in that way but makes sense to me.

1) if energy input required to hit activation energy >> energy difference between reactant and product, you'd get an endothermic reaction
(i.e. invest more energy than you eventually get out)
2) if energy input required to hit activation energy << energy difference between reactant and product, you'd get an exothermic reaction
(i.e. get more energy back than you invested)

graphing both examples should demonstrate that is true
1) big jump to activation energy and we have less of a drop down to products = higher enthalpy of products than reactants so overall endothermic
2) small jump to activation energy and larger drop down to products = lower enthalpy of products than reactants so overall exothermic

it is easier just to compare the total enthalpy of the products vs. the reactants though.