# Calculating the Enthalpy Change of Reaction for the Displacement Reaction Between Zinc and Aqueous Copper Sulphate

Topics: Thermodynamics, Measurement, Energy Pages: 4 (1087 words) Published: April 25, 2011
Calculating the Enthalpy Change of Reaction for the Displacement Reaction between Zinc and Aqueous Copper Sulphate

Data Collection and Processing

Observations:
* Drops of water left on the inside of the measuring cylinder * Hole in the lid, possible escape route for gas or splash-back * The polystyrene cup felt warm during the reaction

By extrapolating the graph we can estimate what the rise in temperature would have been if the reaction had taken place instantaneously. I can conclude that if the reaction had taken place instantaneously, the solution would have reached a temperature of 450C. * Calculate mass of the copper sulphate solution:

mass = volume x density = 25 x 1 = 25 g
* Calculate change in temperature when the zinc is added: ΔT = Tmaximum - Tminimum = 45 – 20 = 25 0C
* Calculate the number moles of copper sulphate reacted:
number of moles = concentration x volume = 0.5 x 0.025 = 0.0125 mol * Calculate energy released during the reaction (Q):
Q = mass of CuSO4 x specific heat capacity x ΔT = 25 x 4.2 x 25 = 2625 J * Calculate ΔH for the reaction (energy change is negative because the reaction was exothermic): ΔH = -Q / number of moles = 2625 / 0.0125 = 210 kJmol-1

‘Absolute Uncertainty’ = abs/unc
‘Percentage Uncertainty’ = %unc

Uncertainty in mass = (%unc in density + %unc in volume) x mass = (0 + 0.5/25) x 25 = ±0.5 g
Uncertainty in temperature change = abs/unc in Tmaximum + abs/unc in Tminimum = 0.5 + 0.5 = ±1 oC
Uncertainty in number of moles = (%unc in concentration + %unc in volume) x number of moles = (0 + 0.5/25) x 0.0125 = 0.00025 = ±0.0003 mol
Uncertainty in energy change = (%unc in mass + %unc in temperature change) x energy change = (0.5/25 + 1/25) x 2625 = 157.5 = ±200 J
Uncertainty in ΔH = (%unc in energy change + %unc in number of moles) x ΔH = (157.5/2625 + 0.00025/0.0125) x 210 = 16.8 = ±20 kJ/mol

Final Value of ΔH: -210 ± 20 kJmol-1
Conclusion
The accepted literature value...

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