How to Calculate Enthalpy Change for a Zinc and Copper Sulfate Reaction?

[mgnymph]

Homework Statement

Zn(solid) + Cu2+(aq) -> Cu(solid) + Zn2+(aqueous)
5.96grams on zinc was added to 25cm^3 of 1M copper(II)sulphate solution in a polystyrene cup.
Over 6 and a half minutes, the temperature of the solution decreased from 68 degrees Celsius to 19.5 degrees C.

1. Calculate enthalpy change for the quantities used, making appropriate assumptions
2. Calculate the enthalpy change for one mole of Zn and CuSO4(aq) and write the thermochemical equation for the reaction

2. The attempt at a solution

1.
moles of Zn = 5.96/65.4 = 0.09113... or 149/1635
therefore, moles of Cu2+ in CuSO4 = 149/1635 (because ratio in above equation is 1 to 1


mass of CuSO4 in solution = 149/1635 x molar mass of CuSO4 = 14.54 grams

q = mcDT,
m = 25 grams, c = 4.2, DT = 68-19.5 = 48.5
q = 5092.5 joules

DH = q/n
= 5092.5/(149/1635)
= 55880Jmol^-1

temperature increased so enthalpy is negative
-> -55880joules per mol

2. Zn + Cu2+ -> Zn2+ + Cu DH = -55880 joules



I'm not sure if I'm doing this right...

 

[Redbelly98]

2. The attempt at a solution

1.
moles of Zn = 5.96/65.4 = 0.09113... or 149/1635
therefore, moles of Cu2+ in CuSO4 = 149/1635 (because ratio in above equation is 1 to 1

For the moles of Cu²⁺, you'll need to use the information given about the copper(II)sulphate solution. It's not necessarily the same # of moles as the Zinc ... either Cu or Zn will be the limiting reactant here.

Other than that, your methods look fine.

 

[Blueshift5]

Your calculations for the enthalpy change appear to be correct. The negative sign indicates that the reaction is exothermic, meaning that it releases energy. This is consistent with the decrease in temperature observed during the reaction.

In terms of the thermochemical equation, it would be written as:

Zn(s) + Cu2+(aq) -> Zn2+(aq) + Cu(s) DH = -55.88 kJ/mol

Note that the enthalpy change is typically reported in kilojoules per mole (kJ/mol) rather than joules per mole (J/mol).

Additionally, the polystyrene cup may have affected the reaction by insulating the solution and preventing heat from escaping. This could have resulted in a slightly lower enthalpy change than the actual value. It is also important to consider any sources of error in the experiment, such as measurement errors or incomplete mixing of the reactants. Overall, the enthalpy change calculated from this experiment may not be entirely accurate, but it can still provide a general understanding of the energy involved in this reaction.