Answer:
The enthalpy change for the solution process [tex]\Delta H_{rxn}[/tex] = - 158.34 kJ/mol
Explanation:
Given that:
The mass of salt AX = 2.499 g
The mass of water = 135.3 g
The mass of the solution = ( 2.499 + 135.3 ) g = 137.799 g
The specific heat of salt solution s is known to be = 4.18 J/g° C
The change in temperature i.e. ΔT = 35.2 °C - 23.6 °C = 11.6 °C
Thus, the amount of heat raised is equal to the heat absorbed by the calorimeter.
∴
[tex]q_{reaction} = q_{solution}[/tex]
[tex]q_{reaction} = -ms_{solution} \Delta T[/tex]
[tex]q_{reaction} = -137.799 \ g \times 4.18 \dfrac{J}{g^0C}\times 11.6^0C[/tex]
[tex]q_{reaction} = - 6682 \ J[/tex]
[tex]q_{reaction} = - 6.682 \ kJ[/tex]
Recall that the mass of the salt = 2.499 g
The number of moles of the salt = [tex]2.499 \ g \times \dfrac{1 \ mol \ of \ AX}{59.1097 \ g}[/tex]
= 0.0422 mol of AX
Finally the enthalpy change, [tex]\Delta H_{rxn} = \dfrac{- 6.682 \ kJ}{ 0.0422 \ mol}[/tex]
= - 158.34 kJ/mol
The enthalpy change for the solution process [tex]\Delta H_{rxn}[/tex] = - 158.34 kJ/mol
