Heat of Solution

The energy change associated with the process in which a solute dissolves in a solvent is called the heat of solution.  This energy change in the net result of two processes, the energy required to break the solute-solute bonds, the crystal lattice energy, and the energy released when the solute particles bond with the solvent molecules, the heat of hydration.  For example, the heat of solution for KCl in water can be considered the sum of the two heat effects.

The crystal lattice energy of KCl, the energy necessary to break apart the KCl crystal lattice and form free ions is represented by equation 1:

[1]                         KCl (s)  à  K¹⁺ (g)    +  Cl^1- (g)         ΔH  =  + 167.6 kcal

The heat of hydration of KCl, the energy released when the free ions are hydrated, is represented by equation 2.

[2]                         K¹⁺ (g)   +  Cl^1- (g)  à  K¹⁺ (aq)   +  Cl^1- (aq)       ΔH = -163.5 kcal

In this example, the overall reaction is endothermic, and the heat of solution is positive since more energy is required in step 1 than is released in step 2.

Overall reaction:  KCl (s)  à  K¹⁺ (aq)   +  Cl^1- (aq)            ΔH  =  + 4.1 kcal

The purpose of this experiment is to determine the heat of solution for some selected compounds from the energy change associated with their solution process.

Procedure:

P1.  Mass a small, clean dry massing pan to the nearest 0.01 g.  Add 1.00 to 1.25 grams of the solid to be studied to the massing pan and re-mass to the nearest 0.01 g.

P2.  Mass a clean, dry calorimeter to the nearest 0.01 g.  Add roughly 15 mL of tap water to the calorimeter and re-mass.  Stir and record the initial temperature of the water in the calorimeter to the nearest 0.1 °C.

P3.  Add the solid to be studied to the calorimeter.  Stir the mixture until the solid is completely dissolved and the temperature remains constant over several readings.  Record the final temperature.

P4.  Rinse and dry both the thermometer and calorimeter, repeating the experiment with the same solid.

P5.  Repeat experiment with the second solid.

Questions and Calculations:

Q1a.  Calculate the mass of the solid, m_solid, by subtracting the mass of the empty massing pan from the mass of massing pan plus solid.

Q1b.  Calculate the number of moles, n, of your assigned solid dissolved in each of your trials by dividing the mass of solid dissolved, m_solid, by its formula mass in g/mol.

Q1c.  Calculate the mass of the water, m_water, by subtracting the mass of the empty calorimeter from the mass of the calorimeter plus water.

Q1d.  Calculate the total mass of the reactants, m_total, by adding the mass of the water, m_water, to the mass of dissolved slat, m_solid.

Q2.  Calculate the change in temperature, Δt, by subtracting the initial temperature of the water,  t_initial, from the final temperature of the solution,  t_final.

Q3a.  Calculate the heat required or liberated when your assigned solid dissolved in water, Δq_water, from the total mass of water plus salt (Q1b), the specific heat capacity for water

(c_p = 1.00 cal/g°C), and the change in temperature (Q2).

Δq_water = m_total • Δt • c_p

Q3b.  Calculate the energy absorbed or released by the calorimeter, Δq_cal, using the calorimeter constant and the change in temperature of the water (Q2)

Δq_cal = calorimeter constant  •  Δt

Q3c.  Calculate the total energy absorbed or released by the solution process, Δq_soln, from the sum of the change in energy of the water (Q3a), plus the change in energy of the calorimeter (Q3b).

Δq_soln  =   Δq_water    +  Δq_cal

Q4a.  Calculate the heat of solution,  ΔH_soln , in cal/mol by dividing the amount of heat liberated or absorbed when the salt dissolves (Q3c) by the number of moles of salt dissolved (Q1a).

ΔH_soln   =  Δq_soln   /  n

Q4b.  Convert each of the heats of reaction in cal/mol into kcal/mol by dividing each ΔH_soln value by 1000 cal/kcal.

Q5.  Calculate the average heat of solution, ΔH_avg, in kcal/mol by averaging the heats of solution (Q4b) for all your trials.

Q6.  Calculate the percent error between your average heat of solution, ΔH_avg, and the theoretical value, ΔH_theor.

% error  =  | ΔH_avg -  ΔH_theor |  •  100

                                                          ________________

                                                                   ΔH_theor

Theoretical heat of solution, ΔH_theor (kcal/mol):NaOH  =  -10.6   exothermic,

          NH₄NO₃  =  6.1   endothermic,       KNO₃     =   8.0  endothermic

Data:  Heat of Solution Lab                                      Name: _____________________

Formula of solid studied: _____________________  Formula mass:  ____________ g/mol

As an ionic compound dissolves in water, the entropy increases.  The ions in the solid phase are arranged in a regular geometric pattern.  For example, sodium chloride crystals have a cubic structure.  In water, however, the ions are separated from one another and are scattered randomly throughout their container giving the solution a much higher entropy.