Bond Enthalpy vs. Standard Enthalpy of Formation

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Labiba Sardar 2A
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Joined: Sat Jul 20, 2019 12:15 am

Bond Enthalpy vs. Standard Enthalpy of Formation

Postby Labiba Sardar 2A » Sun Jan 26, 2020 10:35 am

What's the difference between between bond enthalpies and standard enthalpies of formation? Bond enthalpies are the amount of energy in a bond, but how is that different from standard enthalpies of formation?

Tyler Angtuaco 1G
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Joined: Wed Sep 11, 2019 12:16 am

Re: Bond Enthalpy vs. Standard Enthalpy of Formation

Postby Tyler Angtuaco 1G » Sun Jan 26, 2020 12:27 pm

Bond enthalpy refers to the energy it takes to break a bond in any particular molecule, whereas standard enthalpy of formation is the enthalpy change produced from the formation of a product from its reactants. Keep in mind the standard enthalpy of formation is the enthalpy change for one mole of a molecule in its standard state at 1 atm and 298K.

Ruby Tang 2J
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Joined: Fri Aug 30, 2019 12:15 am

Re: Bond Enthalpy vs. Standard Enthalpy of Formation

Postby Ruby Tang 2J » Sun Jan 26, 2020 1:20 pm

1. Bond enthalpy is the amount of energy required to break one mole of the given bond (ex. if the bond enthalpy of an O-H bond is 464 kJ/mol, then it would take 464 kJ to break 6.022*10^23 O-H bonds).
2. Standard enthalpy of formation of a substance is defined as the change in enthalpy when 1 mole of the substance is formed from its constituent elements in their standard states (ex. CO2(g) formed from C(s) and O2(g)).
3. Both can be used to calculate reaction enthalpy: one can use knowledge of the standard enthalpy of formation of every molecule involved in the reaction in order to solve for the reaction enthalpy. However, when we do not have this information, we can use bond enthalpies instead. With this method, we would figure out which bonds (and how many of those bonds) were broken and which were formed over the course of the reaction.
4. Using bond enthalpies is not as accurate as using standard enthalpy of formation, because bond enthalpies are generally averages of a given bond in a lot of different molecules. For example, the C-H bonds in CH4 vs C2H4 are not the same: one might be slightly stronger and therefore have a higher bond enthalpy. However, we treat them as the same when we use bond enthalpies to calculate reaction enthalpy. The only exception is diatomic molecules, for which the bond enthalpies are experimentally determined.


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