Question

When 2.499 g of AX (s) dissolves in 135.3 g of water in a coffee-cup calorimeter the temperature rises from 23.6 °C to 35.2 °C. Calculate the enthalpy change (in kJ/mol) for the solution process. AX space (s )space rightwards arrow straight A to the power of plus (a q )space plus space straight X to the power of minus space (a q )Assumptions for this calculation: The specific heat of the solution is the same as that of pure water (4.18 J/gK) The density of water = 1.000 g/mL The liquid’s final volume is not changed by adding the solid The calorimeter loses only a negligible quantity of heat. The formula weight of AX = 59.1097 g/mol. Be sure you include the correct sign for the enthalpy change.

Answer 1

Answer:

The enthalpy change for the solution process $\Delta H_{rxn}$ = - 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.

∴

$q_{reaction} = q_{solution}$

$q_{reaction} = -ms_{solution} \Delta T$

$q_{reaction} = -137.799 \ g \times 4.18 \dfrac{J}{g^0C}\times 11.6^0C$

$q_{reaction} = - 6682 \ J$

$q_{reaction} = - 6.682 \ kJ$

Recall that the mass of the salt = 2.499 g

The number of moles of the salt = $2.499 \ g \times \dfrac{1 \ mol \ of \ AX}{59.1097 \ g}$

= 0.0422 mol of AX

Finally the enthalpy change, $\Delta H_{rxn} = \dfrac{- 6.682 \ kJ}{ 0.0422 \ mol}$

= - 158.34 kJ/mol

The enthalpy change for the solution process $\Delta H_{rxn}$ = - 158.34 kJ/mol