The end goal of this problem is to find the enthalpy of solution for CaCl2. It\'

Question

The end goal of this problem is to find the enthalpy of solution for CaCl2. It's given in a table as -81.3 (kj/mole)

So, based on experimental data:

I have 5.1g (0.0460 moles) of CaCl2. Dissolved in 50g water, and the water temperature change is 23.7 to 32.0 (-8.3) degrees.

Total Heat (Q) Transferred upon dissolving (J), I've calculated as such: (50g+5.1g) * 4.184 * (-8.3) = -1913.47 J

So, to get the enthalpy of solution, then you take kJ/mole: -1.913 kJ / 0.0460 moles = -41.59.   

But this isn't even close to -81.3. I can't figure out what has gone wrong.

Explanation / Answer

Answer:

q = mass(water) x specific heat capacity(water) x change in temperature(solution)
q = mH2O(l) x cg x (Tf - Ti)
q = 50 x 4.184 x (-8.3) = -1736.36 J

2. Calculate the enthalpy change, ?H, in kJ mol-1 of solute:

?H = -q/1000 ÷ n(solute) = -1736.36/1000 ÷ 0.0460 = -37.73 kJ mol-1
?H is negative because the reaction is exothermic (energy is released causing the temperature of the solution to increase).

this is the way to solve the problem.

Dear student, your way of solving the problem is correct but the values given in the data are wrong( either moles/water quantity/rise in temperature). thats way you can't figure out the value which is given in table (-81.3Kj).

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