4 Calorimetry

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Calorimetry Measurement of Enthalpy Change

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Arizona State University 1 School of Molecular Sciences How Can the Thermodynamics of Dissolution be Defined? Introduction: The Gibbs-Helmholtz equation expresses the relationship between the free-energy change, the enthalpy change, and the entropy change at constant temperature and pressure: ∆G = ∆H - T∆S Equation 1 From knowing the value of ∆G, you may predict whether a process/reaction will be spontaneous at a certain temperature. A process is spontaneous if ∆G is negative (∆G < 0), nonspontaneous if ∆G is positive (∆G > 0), and at equilibrium if ∆G = 0. The enthalpy change, ∆Hrxn, 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 ∆H) because energy is needed to break bonds, and energy released (exothermic, negative ∆H) because energy is released as new bonds are formed. ∆Hrxn 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 ∆Grxn to check your prediction. From the law of conservation of energy (energy is conserved) the total energy for the dissolution process is: qsystem + qsurroundings = 0 or qsystem = - qsurroudnings Equation 2 where qsystem (or qrxn) represents the heat gained or lost by dissolving the solid (the system), and qsurroundings (or qsolution) 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 = (mass solution)*(specific heat of solution)*(change in temperature) q = m· Cs · T, where T=Tfinal-Tinitial Equation 3 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 is1.00 g/mL.) Because the solution is very Arizona State University 2 School of Molecular Sciences dilute, the specific heat of the solution is basically equal to that of water, which is 4.184 J/gºC. To calculate change (∆) for a variable it is always final minus initial. The value of qsurr can be calculated.

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