This is probably insanely easy, but it just doesn't seem right to me. The question is:
Zinc dissolves in acid according to the balanced chemical reaction:
Zn(s) + 2 H+(aq) -> Zn2+(aq) + H2(g)
A sample of zinc is placed in the ice calorimeter described in the "Experimental" section (I'll specify below). If 0.0657g of zinc causes a decrease of 0.109mL in the ice/water volume of the calorimeter, what is the enthalpy change, per mole of zinc, for the above reaction per mole of zinc.
The calorimeter details:
T = constant
density of water = 1.000 g mL-1
density of ice = 0.917 g mL-1
H2O (s) -> H2O (l) deltaH of fusion = 6.01 kJ mol-1, or 333 J g-1
It can be determined that 3.68 kJ are released per mL change in the volume of the ice/water mixture.
The attempted answer:
So qrxn = -qcalorim
temperature is constant, so T isn't needed
pressure is constant, so q = deltaH
Since we're just finding the change in enthalpy per mole of zinc, deltaH = qzinc
qzinc = mc = (0.0657g)(0.390 J g-1) = 0.026 J
. doesnt seem right.
UPDATE:
Amount of moles for Zn = 0.0657g / 65.39 g mol-1 = 0.001 mol
0.109mL x 3.68 kJ mL-1 = 0.401 kJ
0.401 kJ / 0.001 mol = 401 kJ mol-1
Is that right?