January 22, 2017

Homework Help: Chemistry

Posted by a Canadian on Wednesday, February 13, 2013 at 5:41pm.

I did a lab today in class and part of the lab is to experimentally find the enthalpy for the dissolution of ammonium chloride by dissolving 4.0 g of NH4Cl(s) in 50.0 mL of H2O(l).

I got a -5°C temperature change and calculated delta H to be 15 kJ/mol. According to the Internet, the actual value is 14.7 kJ/mol? So I'm a little over. My initial worry was that For sources of error, we only ever discussed things that could cause heat loss - not anything that would actually result in a higher heat gain. And "maybe we measured wrong" definitely wouldn't work as a legitimate source of error. But I thought about it a bit more and the positive 15 kJ/mol indicates that the reactants absorbed that much energy right? Which would result in a lower temperature. And heat loss through the calorimeter, etc would make the temperature lower and make it seem like the reactants absorbed more energy than they really did, and so the enthalpy change was calculated to be a higher value. I was wondering if this is right?

I hope I made sense.

It's also very possible that the Internet/my calculations wrong.

Also, the other part of my lab was to find the value of delta H for the reaction between 25 mL of 1.0M HCl(aq) and 25 mL of 1.0 M NH3(aq). I tried looking for the correct value on the Internet to check if my value, -41.84 kJ/mol, was close and to make sure I didn't completely mess up my lab, but I couldn't find it. I was wondering if anyone just happened to know..?

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