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Chemistry Forums for Students => Undergraduate General Chemistry Forum => Topic started by: nortorius on October 26, 2006, 10:03:24 PM

Title: Calorimeter prelab exercise
Post by: nortorius on October 26, 2006, 10:03:24 PM

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 q_rxn = -q_calorim

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 = q_zinc

q_zinc = 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?