Hess's Law - Thermochemical Equations

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In summary: What do you think is going on?In summary, the problem is that the enthapy of formation for NH4CL is different from the accepted standard enthalpy of formation. I don't know what the accepted value is, but it might be something that I am missing.
  • #1
kevinzak
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Hello! I am a high school senior taking an advanced chemistry course. I have been working on this problem for two days straight, and cannot figure out our error. Here is the problem we are given:

Calculate the standard enthalpy of formation for the reaction HCL (g) + NH3 (g) ---> NH4CL (s), given the following thermochemical equations:

H2 (g) + Cl2 (g) ---> 2HCL (g), standard enthalpy of formation: -184kJ (-92kJ/mol)
N2 (g) + 3H2 (g) ---> 2NH3 (g), standard enthalpy of formation: -92kJ (-46kJ/mol)
N2 (g) + 4H2 (g) + Cl2 (g) ---> 2NH4CL (s), standard enthalpy of formation: -628kJ (-314kJ/mol)

It wants us to combine the equations using Hess's Law to arrive at the desired equation, and thus, the desired enthalpy of formation. I did so, and I arrive at the desired equation given above ( HCL (g) + NH3 (g) ---> NH4CL (s) ). However, the number I have arrived at time and time again (-176kJ/mol) is not the accepted standard enthalpy of formation for NH4CL, -314.43kJ/mol.

Someone please show me the error before I bash my head against the wall any further! Thanks!

Note, the book that gives the problem lists the answer I get (-176kJ/mol) as correct. I am wondering why this is not the accepted value. I'm sure it is something simple that I am missing, but I have been staring at it for much too long now and must admit I need assistance.
 
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  • #2
You are mixing up enthapy of formation with enthapy of reaction.

The chemical equation that you have does not define the enthapy of formation; what you have found is the enthalpy change associated with that particular reaction.

The "enthapy of formation" of a particular compound is defined as the energy change associated with the production of that compound from its constituent elements in their naturally occurring form.

For example, for ammonium chloride, the enthapy of formation is the energy change associated with:

[tex]\frac{1}{2}{N_{2(g)}} + 2{H_{2(g)}} + \frac{1}{2}C{l_{2(g)}}\rightarrow N{H_4}C{l_{(s)}}[/tex]
 
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  • #3
Hmm, then my textbook is simply incorrect in it's wording of the problem? I copied the problem verbatim, and it definitely says find the enthalpy of formation for that reaction.
 

1. What is Hess's Law and how is it related to thermochemical equations?

Hess's Law is a principle in chemistry that states the total enthalpy change of a reaction is independent of the pathway taken from reactants to products. It is related to thermochemical equations because it allows us to calculate the enthalpy change of a reaction by using known enthalpy values of other reactions.

2. How do you use Hess's Law to calculate the enthalpy change of a reaction?

To use Hess's Law, you need to have a series of reactions that can be added together to give the overall reaction you are interested in. Then, you can use the enthalpy values of each reaction to calculate the enthalpy change of the overall reaction by applying the law of conservation of energy.

3. Can Hess's Law be applied to all types of reactions?

Yes, Hess's Law can be applied to all types of reactions, including both exothermic and endothermic reactions.

4. What is the significance of Hess's Law in thermochemistry?

Hess's Law is significant in thermochemistry because it allows us to determine the enthalpy change of a reaction without directly measuring it. This is especially useful in cases where the reaction is difficult to carry out or measure directly.

5. Are there any limitations to using Hess's Law?

One limitation of using Hess's Law is that it assumes the reactants and products are in the same physical state in all reactions. It also does not account for changes in temperature or pressure, which can affect the enthalpy of a reaction. Additionally, it requires accurate and precise enthalpy values for each reaction, which may not always be available.

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