Answer to Question #119311 in Organic Chemistry for emily

Question

A student is determining the enthalpy of solution for ammonium nitrate by adding ammonium nitrate to a calorimeter and measuring the temperature change. The accepted value is ΔHsol’n = 25.7 kJ/mol. A student, not paying attention, adds 1.5g of ammonium nitrate instead of 1.0g. When they perform their calculations, they use 1.0g as the mass. Will their calculation result in an answer that is higher than the accepted value or lower? Explain. (2 marks)

Answer 1

Dissolution of ammoinum nitrate is endothermic (positive enthalpy value), so temperature decreases. In real life he added more salt than should be added, so the registered temperature change will be larger. When he calculates the experimental enthalpy change normalized per 1.0 g, he will get higher enthapy than the accepted one. Because to get the correct enthaply he should had normalized it per 1.5 g ammonium nitrate.