Instant cold packs used to treat athletic injuries contain solid NH4NO3 and a pouch of water. When the pack is squeezed, the pouch breaks and the solid dissolves, lowering the temperature because of the following endothermic reaction.

NH4NO3(s) + H2O(l) = NH4NO3(aq) deltaH = +25.7 kJ

 

What is the final temperature in a squeezed cold pack that contains 50.0 g of NH4NO3 dissolved in 115 mL of water? Assume a specific heat of 4.18 J/(g·°C) for the solution, an initial temperature of 25.0°C, and no heat transfer between the cold pack and the environment. To find the mass of water use the density of water = 1.0 g/mL. Hint: The process takes place at constant pressure.

 

So i know at constant pressure i have to use

 

Cs * M of solution * deltaT = -(molesNH4NO3)*DeltaH of reaction

 

Can someone please set this problem up because I tried to set it up but had difficulties with some of the values.

Edited February 1, 201114 yr by Don H.

You have every value needed to find T2, which is T1 + deltaT. This is just a fill the numbers-exercise.

You will also need to find out how many moles of NH4NO3 you are dissolving.

He has the mass of the ammonium nitrate. Not sure if the solubility of that stuff is 50g/115ml tho, you might want to check that