A 'QwikCure' pack, used to treat sporting injuries, contains a bag of water inside a larger bag of finely powdered ammonium nitrate, NH4NO3 - VCE - SSCE Chemistry - Question 2 - 2009 - Paper 1

For this part, you should create an energy profile diagram that illustrates the energy changes during the dissolution process of NH4NO3. The diagram should depict an initial energy level for NH4NO3(s), a higher peak representing the transition state, and a lower energy level for the resulting NH4NO3(aq). Label the axes appropriately: energy (kJ mol⁻¹) on the y-axis and reaction progress on the x-axis.

The equilibrium reaction can be written as:

$NH_4^+(aq) + H_2O(l) \rightleftharpoons NH_3(aq) + H_3O^+(aq)$

The change that occurred at time t2 was likely due to a decrease in concentration of either ammonium ions or water. When the concentration of the reactants decreases, the rate of the forward reaction also decreases. Consequently, the rate of the back reaction will initially increase due to the shift in equilibrium as the system adjusts to restore balance.

The value of the equilibrium constant at time t2 would be equal to the value at time t1, as equilibrium constants remain unchanged unless temperature changes.

The expression for the acidity constant K_a for the reaction is given by:

$K_a = \frac{[NH_3][H_3O^+]}{[NH_4^+]}$

The pH of the solution is given as 5.04. To find the concentration of H3O⁺ ions, we can use the formula:

$[H_3O^+] = 10^{-pH} = 10^{-5.04} \approx 9.12 \times 10^{-6} \text{ mol L}^{-1}$

First, we need to calculate the moles of NH4NO3 dissolved in 300 mL of the solution:

Using the formula:

$ext{moles} = [NH_4NO_3] \times ext{volume}$

Given that the pH 5.04 indicates the presence of H3O⁺ ions, we can assume initial concentration matches that of NH4NO3.

Using the molarity, calculate the mass as follows:

$ext{Mass} = ext{moles} \times ext{molar mass}$
Assuming the molar mass of NH4NO3 is approximately 80.04 g/mol, apply moles to find the mass.