1. The problem statement, all variables and given/known data 20g of NH4NO3 (MM=80.05 g/mol) if dissolved in 125g of water in a coffee-cup colorimeter, the temperature falls from 23.5C to 13.4C. Determine 'q' for the dissolving of the compound. Is the process exothermic or endothermic? 2. Relevant equations q(calorimeter)=C(calorimeter)*[Delta]T 3. The attempt at a solution [Delta]T= -10.1 C(calorimeter)=~n(water)*C(water) = 125g H2O/18.016g/mol = 6.9382mol H2O Ccal=~6.9382mol H2O* 75.291 J/mol C = 522.4 J/C qcal=52.4J/C*-10.1C= -5.28kJ q(calorimeter)= -q(chemical) = +5.28kJ This answer makes sense, however, I don't understand why I am not taking into account the amount of moles NH4NO3. A larger amount of moles would make a larger deviation in T, which would change q. Why are moles NH4NO3 not taken into account? Also, I'm using the heat capacity (C) of water, not of NH4NO3. The question is asking for 'q' for the dissolving compound. Is this correct?