Energy Exchange in N2 + H2 Reaction: Examining the Bond Changes

[Erin_Sharpe]

When N2 and H2 are reacted in a bomb calorimeter to form NH3 the temperature of the bomb calorimeter and the water it contains increases.
Is the energy required to break the bonds in the reactants greater or less than the energy released during the formation of the bonds of the prouducts.

I said it was less...but I don't think that's right. Can someone help?

Thanks,
Erin

 

[maverick280857]

You have to reason based on experimental evidence but generally nitrogen is considered to have a high bond dissociation enthalpy so you need to invest quite a large amount of energy per mole to break nitrogen-nitrogen bonds in N2 gas. At the same time if you are told that heat is released all that you can say is that delta(H) < 0 or equivalently

$$\Delta H_{product} < \Delta H_{reactants}$$

Cheers
Vivek

 

[ravenprp]

Hi Erin,

You were correct in saying that the energy required to break the bonds in the reactants is less than the energy released during the formation of the bonds in the products. This is due to the fact that the formation of bonds releases energy, while breaking bonds requires energy.

In the reaction between N2 and H2 to form NH3, the N-N triple bond and the H-H single bonds are broken in the reactants. These bonds require a certain amount of energy to break, called the bond energy. However, when the N-H bonds are formed in the product NH3, energy is released. This energy released is greater than the energy required to break the bonds in the reactants, resulting in a net release of energy.

This excess energy is what causes the temperature of the bomb calorimeter and the water to increase. Therefore, the energy released during the formation of the bonds in the products is greater than the energy required to break the bonds in the reactants.

I hope this helps clarify your understanding. Let me know if you have any other questions.