Question

A student was required to carry out an experiment in the lab to determine the enthalpy of solution of ammonium nitrate. The student was given a polystyrene cup, a measuring cylinder and a thermometer

A. Why was a polystyrene cup and not a beaker used in the experiment? [1]

B. State one assumption that was made when performing the calculations for this reaction. [1]
C. Given that the initial temperature of the water was 27.3 C and the final was 15.1C state whether the reaction was endothermic or exothermic? [1]

d. Using the above information, calculate the enthalpy of the solution in kJ/mol when 16.0g of ammonium nitrate dissolves in 100cm3 of water. [2] (Give answer to 3 significant figures

Answer 1

A student was required to carry out an experiment in the lab to determine the enthalpy of solution of ammonium nitrate. A polystyrene cup was used because polystyrene is a heat insulator that would prevent heat loss to the surroundings. In that way, we can assume that heat loss is negligible in the endothermic reaction. The enthalpy of the solution is 29.6 kJ/mol.

A student was required to carry out an experiment in the lab to determine the enthalpy of solution of ammonium nitrate. The student was given a polystyrene cup, a measuring cylinder and a thermometer.

A. Why was a polystyrene cup and not a beaker used in the experiment?

Because polystyrene is a heat insulator that would prevent heat loss to the surroundings.

B. State one assumption that was made when performing the calculations for this reaction.

The heat loss to the surroundings is negligible.

C. Given that the initial temperature of the water was 27.3 °C and the final was 15.1 °C, state whether the reaction was endothermic or exothermic?

Since the final temperature was lower than the initial temperature, the reaction absorbed heat, which means it is endothermic.

D. Using the above information, calculate the enthalpy of the solution in kJ/mol when 16.0 g of ammonium nitrate dissolves in 100 cm³ of water. (Give answer to 3 significant figures)

The molar mass of ammonium nitrate is 80.04 g/mol. The moles corresponding to 16.0 g of ammonium nitrate are:

$n = 16.0 g \times \frac{1mol}{80.04g} = 0.200 mol$

We have 16.0 g of solute and 100 cm³ of water (100 g of water assuming density 1 g/cm³). The mass of the solution is:

$m = 16.0 g + 100 g = 116 g$

We can calculate the heat absorbed (Q) by the water using the following expression.

$Q = c \times m \times \Delta T$

where,

• c: specific heat of the solution (we will assume it is equal to that of water)
• m: mass of the solution
• ΔT: change in the temperature

$Q = \frac{4.184J}{g. \° C } \times 116 g \times (15.1 \° C - 27.3 \° C) \times \frac{1kJ}{1000J} = -5.92 kJ$

According to the law of conservation of energy, the sum of the heat absorbed by the water (Qw) and the heat released by the solution (Qs) is zero.

$Qw + Qs = 0\\Qs = -Qw = 5.92 kJ$

Finally, we can calculate the enthalpy of the solution (ΔH) using the following expression.

$\Delta H = \frac{Qs}{n} = \frac{5.92kJ}{0.200mol} = 29.6 kJ/mol$

A student was required to carry out an experiment in the lab to determine the enthalpy of solution of ammonium nitrate. A polystyrene cup was used because polystyrene is a heat insulator that would prevent heat loss to the surroundings. In that way, we can assume that heat loss is negligible in the endothermic reaction. The enthalpy of the solution is 29.6 kJ/mol.

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