Question

A calorimeter contains 35.0 mL of water at 11.5 ∘C . When 1.30 g of X (a substance with a molar mass of 66.0 g/mol ) is added, it dissolves via the reaction X(s)+H2O(l)→X(aq) and the temperature of the solution increases to 29.5 ∘C . Calculate the enthalpy change, ΔH, for this reaction per mole of X. Assume that the specific heat of the resulting solution is equal to that of water [4.18 J/(g⋅∘C)], that density of water is 1.00 g/mL, and that no heat is lost to the calorimeter itself, nor to the surroundings. Express the change in enthalpy in kilojoules per mole to three significant figures.

Answer 1

First of all, we need to calculate the total energy change using the formula Q=mcθ, where m is the mass of water, c is the specific heat capacity of water and θ is the temperature change.

Q= (35)(4.18)(29.5-11.5)
= 2633.4J

Note: The mass of water is 35g since its density is 1g/ml.

We then need to determine the no of moles of X;
66g= 1 mole of X
1.3g= 1.3/66
= 0.0197
The dissolution of 0.0197 moles of X produces 2633.4J of heat energy
For 1 mole: 2633.4/0.0197
=133000J
= 133KJ
Thus, the enthalpy change,ΔH= 133kJ/mol