Thermochemistry: Calorimetry and Hess’s Law Some chemical reactions are endothermic and proceed with absorption of heat while others are exothermic and proceed with an evolution of heat. The magnitude of the heat change is determined by the particular reaction of interest, as well as by the amount of reactants consumed. The change in enthalpy, ∆H, is the heat (qp) associated with a reaction atconstant pressure, in which no work is involved other than "expansion work" or "compression work", associated with volume changes of the system (PV work): qp = ∆H The value of ∆H is a characteristic and important property of the reaction; convenient sets of units for the enthalpy is kJ/(mole of a component of the reaction). These units emphasize that enthalpy is an extensive property. Enthalpies ofreaction (often called "heats of reaction") are usually tabulated for reactions occurring isothermally (at constant temperature), usually 25 °C. As long as the temperature change is not very large, the assumption that ∆H is constant with T is reasonable and within experimental error. In this experiment, you will use an adiabatic, solution calorimeter, to determine enthalpies of reaction. TheAdiabatic Calorimeter An adiabatic calorimeter is a device so well thermally insulated that the transfer heat across the boundary between the calorimeter and its surroundings is effectively zero during the experiment. In practice, perfect insulation is never achieved, but the error can be minimized through proper data treatment. In part I of this experiment, the adiabatic calorimeter consists of adouble Styrofoam cup fitted with a thermometer calibrated in tenths of a degree (permitting temperatures to be estimated to about ±0.02°C) and a single, inverted Styrofoam cup (with a small hole for the thermometer) as a lid. It is crude, but very effective because of the excellent insulating properties of Styrofoam. The heat exchanged between the reaction (∆H) and the surroundings (adiabaticcalorimeter) can be determined from: ∆H = – Ccalorimeter∆Tcalorimeter where heat capacity, Ccalorimeter, is the heat (of reaction) required to raise the temperature of the calorimeter by one degree. It will have contributions from the thermometer, the cups, the lid, and, most importantly, the water (solvent) in the calorimeter. In fact, we will assume that the contribution of the water in the calorimetergreatly outweighs all other contributions. The equation can then be written as: ∆H = – Cwater∆Tcalorimeter
Since the specific heat capacity, s, of water (4.184 J/g·oC) is a well know quantity, all that is required to obtain the heat capacity of an aqueous solution calorimeter is to determine the (approximate) mass of water in it and use ∆H = – mwaterswater∆Tcalorimeter Key techniques forobtaining accurate results are starting with a dry calorimeter, measuring solution volumes precisely, and determining ∆T accurately. Careful experimenters deal with the first two items easily. The last is somewhat more difficult. The change in temperature is determined by measuring the initial temperature, T1, of the reactants, and the maximum (or minimum) temperature, T2, of the contents of thecalorimeter during the exothermic (endothermic) reaction. The determination of a precise value for T2 is complicated by the fact that a small heat exchange occurs between the surroundings and the contents of the calorimeter, both during the reaction and after its completion. The exact rate of exchange depends on the insulating properties of the calorimeter. A correction for this heat loss is made by anextrapolation of a temperature vs. time curve before and after the reaction (Figure 1).
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Figure 1: Temperature versus Time Curve. From extrapolation: Ti = 24.0 oC and Tf = 34.8oC. T1 and T2 are determined by extrapolating the linear portion of each curve forward or back to zero time (the...
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