5.1: Background

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All chemical reactions involve either the evolution or absorption of energy, usually in the form of heat. Heat is released in exothermic reactions and absorbed in endothermic reactions. The change in energy can be shown in an energy diagram. An energy diagram for the combustion of solid carbon is shown below.

Since this reaction is endothermic, the system absorbs energy from the surroundings and the energy of the products that remain is greater than the energy of the original reactants. ΔE for an endothermic reaction is positive indicating that the energy of the system has increased. For exothermic reactions, energy is released by the system, and the energy of the products is less than the energy of the reactants. ΔE is negative indicating that the energy of the system has decreased.

The energy change that occurs during a chemical reaction can be determined by measuring the temperature change that occurs when the reaction is carried out in an insulated container known as a calorimeter. The amount of heat involved depends upon the conditions under which the reaction is conducted. For example, most chemical reactions are carried out at constant pressure, usually atmospheric pressure. The heat evolved or absorbed at constant pressure is referred to as the enthalpy change, ΔH.

The temperature change experienced by a substance depends upon its heat capacity, C, which is defined by the equation

(1)

where Q is the quantity of heat added to the substance, T_i is the initial temperature of the substance, T_f is the final temperature of the substance, and DT is the change in temperature of the substance. Note that heat capacity and the temperature change are inversely related; that is, the greater the heat capacity, the smaller the temperature change for a given quantity of heat.

Another quantity related to heat capacity is the specific heat capacity, c. Some books use the symbol, C_s, for specific heat capacity. The specific heat capacity is the heat capacity per unit mass, m, of a substance and is defined as

(2)

The specific heat capacity of water is 4.184 J/^oCg, which means that 4.184 J of heat are required to increase the temperature of 1 gram of water by 1 ^oC. The heat capacity of aluminum is 0.89 J/^oCg, thus only 0.89 J of heat are needed to raise the temperature of 1 gram of aluminum by 1 ^oC.

To determine the quantity of heat required to increase the temperature of 10.0 grams of water from 25 ^oC to 35 ^oC, rearrange equation (2) to the following form.

Q = m c ΔT (3)

Q = (10.0 g)(4.184 )(10 ^oC) = 4.2 x 10² J

Note that the temperature change has only two significant figures; therefore, the final answer is limited to only two significant figures.

In this experiment you will measure the heat involved in the reaction of a strong acid and a strong base. This enthalpy change is called the heat of neutralization. The heat given off in this reaction is mainly absorbed by the resulting solution; however, a small amount of heat is absorbed by the calorimeter.

heat given off by reaction = heat absorbed by solution + heat absorbed by calorimeter cup

or ΔH = Q_solution + Q_calorimeter

ΔH = m_solution x c_solution x DT + m_calorimeter_cup x c_calorimeter_cup x ΔT

The quantity (m_calorimeter_cup x c_calorimeter_cup) is the heat capacity of the calorimeter cup, C_calorimeter_cup, and is known as the calorimeter constant. The calorimeter constant is determined by measuring the temperature change for a process for which the quantity of heat is accurately known. For the styrofoam cups used for the calorimeter in this experiment, the calorimeter constant is 20. J/^oC, that is, for each 1 ^oC increase in temperature the calorimeter absorbs 20. J of heat.

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