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Calorimetry for AP Chemistry

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By — McGraw-Hill Professional
Updated on Feb 9, 2011

Practice problems for these concepts can be found at:

Calorimetry is the laboratory technique used to measure the heat released or absorbed during a chemical or physical change. The quantity of heat absorbed or released during a chemical or physical change is represented as q and is proportional to the change in temperature of the system being studied. This system has what is called a heat capacity (Cp), which is the quantity of heat needed to change the temperature 1 K. It has the form:

    Cp = heat capacity = qT

Heat capacity most commonly has units of J/K. The specific heat capacity (or specific heat) (c) is the quantity of heat needed to raise the temperature of 1 g of a substance 1 K:

    c = q/(m × ΔT ) or q = cmΔT,

where m is the mass of the substance.

The specific heat capacity commonly has units of J/g.K. Because of the original definition of the calorie, the specific heat capacity of water is 4.184 J/g.K. If the specific heat capacity, the mass, and the change of temperature are all known, the amount of energy absorbed can easily be calculated

Another related quantity is the molar heat capacity (C), the amount of heat needed to change the temperature of 1 mol of a substance by 1 K

Calorimetry involves the use of a laboratory instrument called a calorimeter. Two types of calorimeter, a simple coffee-cup calorimeter and a more sophisticated bomb calorimeter, are shown in Figure 9.1. In both, a process is carried out with known amounts of substances and the change in temperature is measured.

The coffee-cup calorimeter can be used to measure the heat changes in reactions or processes that are open to the atmosphere: qp, constant-pressure reactions. These might be reactions that occur in open beakers and the like. This type of calorimeter is also commonly used to measure the specific heats of solids. A known mass of solid is heated to a certain temperature and then is added to the calorimeter containing a known mass of water at a known temperature. The final temperarure is then measured allowing us to calculate the ΔT. We know that the heat lost by the solid (the system) is equal to the heat gained by the surroundings (the water and calorimeter, although for simple coffee-cup calorimetry the heat gained by the calorimeter is small and is ignored)

    qsolid = qwater

Calorimetry

Substituting the mathematical relationship for q gives:

    –(csolid × msolid × ΔTsolid) = cwater × mwater × ΔTwater

This equation can then be solved for the specific heat capacity of the solid.

The constant-volume bomb calorimeter is used to measure the energy changes that occur during combustion reactions. A weighed sample of the substance being investigated is placed in the calorimeter, and compressed oxygen is added. The sample is ignited by a hot wire, and the temperature change of the calorimeter and a known mass of water is measured. The heat capacity of the calorimeter/water system is sometimes known.

For example, a 1.5886 g sample of glucose (C6H12O6) was ignited in a bomb calorimeter. The temperature increased by 3.682°C. The heat capacity of the calorimeter was 3.562kJ/°C, and the calorimeter contained 1.000 kg of water. Find the molar heat of reaction (i.e., kJ/mole) for

Practice problems for these concepts can be found at:

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