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# Constant-Pressure Calorimetry

CHEM-LL8BGB

Consider the following reaction: $N{H}_{4}Cl _{(s)} \rightarrow {N{H}_{4}}^{+} _{(aq)} + {Cl}^{-} _{(aq)}$

$3.00\ g$ of $N{H}_{4}Cl$ are dissolved in $40.0\ g$ of water in a coffee-cup (constant-pressure) calorimeter. The initial temperature of the water is $22.0\ ˚C$. After the $N{H}_{4}Cl$ dissolves, the temperature is $17.2\ ˚C$.

For this problem, ignore the heat capacity of the coffee-cup calorimeter. Use $4.18\ \frac{J}{g \cdot ˚C}$ as the heat capacity for water. Assume no heat is lost to the surroundings.

Calculate $\Delta H$ for the dissolution reaction above in $\frac{kJ}{mol}$.

A periodic table is provided for reference: