Reference no: EM131220684

How Can the Thermodynamics of Dissolution be Defined?

Introduction:

The Gibbs-Helmholtz equation expresses the relationship between the free-energy change, ΔG, the enthalpy change, ΔH, and the entropy change, ΔS, at constant temperature and pressure:

ΔG = ΔH - TΔS .......................................... Equation 1

From knowing the value of AG, you may predict whether a process/reaction will be spontaneous at a certain temperature. A process is spontaneous if ΔG is negative (ΔG a 0), nonspontaneous if ΔG is positive (ΔG > 0), and at equilibrium if ΔG = 0.

The enthalpy change, ΔH_rnx is the heat gained or lost by a system during a reaction carried out at constant pressure. Most reactions occur in several steps, with energy required (endothermic, positive ΔH) to break bonds, and energy released (exothermic, negative ΔH) released as new bonds are formed. ΔH_rnx represents the total change in heat energy or enthalpy over the course of the reaction.

In this experiment, you will use a coffee-cup calorimeter to determine the heat absorbed or released during the dissolution of ammonium chloride and the dissolution of calcium chloride. From observing the contents of the coffee-cup calorimeter, you will decide whether the dissolution processes are spontaneous or nonspontaneous. You will also calculate values of ΔG_rnx to check your prediction.

From the law of conservation of energy (energy is conserved) the total energy for the dissolution process is:

q_system + q_{surrounclings} = 0 or q_system = - q_surroudnings ....................... Equation 2

where q_system (or q_rnx) represents the heat gained or lost by dissolving the solid (the system), and q_{surrounclings} (or q_solution) is the heat gained or lost by the solution in the calorimeter (the surroundings). Thus, heat energy is essentially transferred between the dissolving solid and the solution in the calorimeter. (For this experiment, the heat absorbed by the cup, probe, and surroundings can be considered as negligible, so it is not included in the expression above.)

The heat absorbed or released by the contents of the calorimeter is given by:

q = m· C_s· ΔT .......................  Equation 3

where m is the mass of the solution, Cu represents the specific heat of solution, and AT is the change in temperature (ΔT=T_final - T_initial)

For this experiment, the mass of the solution is the sum of the masses of the water and solid placed in the calorimeter. (Recall that the density of water is 1.00 g/mL). The specific heat of the solution can be assumed to be equal to that of water because the solution is very dilute. C_s for water is 4.184 J/g°C. The value of q_surr can be calculated by plugging in experimental values into Equation 3, and then determining the heat of reaction, q_sys, from Equation 2.

The molar enthalpy of reaction, ΔH_rnx, will then be calculated by dividing the heat of reaction by the experimental number of moles of salt used in the experiment.

ΔH_rnx., = q_sys, / moles salt         ..............................Equation 4

You will need to calculate the ΔS°_rnx. values for the dissolution of solid ammonium chloride and calcium chloride using data from Appendix C in the back of your textbook. We do not have experimental data for this calculation, so we will use the textbook values and solve for ΔS°_rnx.

Δ S^o_rnx = Σn5^o(products)- ΣmS°(reactants) .................. Equation 5

Finally you can calculate the experimental change in Gibb's Free Energy (ΔG) using the Gibbs-Helmholtz equation, Equation 1, using the initial temperature for T, the experimental value for the enthalpy of reaction, and the textbook value for the entropy of reaction.

o Introduce the dissolution. in water general reaction.

o What is thermodynamics?

o What differences/similarities did you see with both solids?

o Explain the importance of taking the mass of each individual substance involved in a system.

o Explain the importance of units for each variable involved in Gibbs Free Energy equation.