0.11 Measuring energy changes in chemical reactions  (Page 5/8)

Calorimetry: Measuring the Heat of a Chemical Reaction

Let’s illustrate by analyzing the example of burning butane instead of methane. If we burn 1.0 g of butane and allow the heat evolved to warm 1.0 kg of water, we observe that an increase in the temperature of the water of 11.8 °C. Therefore, by [link] , elevating the temperature of 1000 g of water by 11.8 °C must require 49,520 J = 49.52 kJ of heat. Therefore, burning 1.0 g of butane gas produces exactly 49.52 kJ of heat.

The method of measuring reaction energies by capturing the heat evolved in a water bath and measuring the temperature rise produced in that water bath is called calorimetry . Following this procedure, we can straightforwardly measure the heat released or absorbed in any easily performed chemical reaction.

By convention, when heat is absorbed during a reaction, we consider the quantity of heat to be a positive number: in chemical terms, q>0 for an endothermic reaction. When heat is evolved, the reaction is exothermic and q<0 by convention.

Observation 3: hess' law of reaction energies

The method of calorimetry we have developed works well provided that the reaction is easily carried out in a way that we can capture the energy transfer in a known quantity of water. But many reactions of great interest are very difficult to carry out under such controlled circumstances. Many biologically important chemical reactions may require the conditions and enzymes only available inside a cell. For example, conversion of glucose C ₆ H ₁₂ O ₆ to lactic acid CH ₃ CHOHCOOH is one of the primary means of providing energy to cells:

Measuring the energy of this reaction is important to understanding the biological process of energy transfer in cells. However, we can’t simply put glucose in a beaker and wait for it to turn into lactic acid. Very specific conditions and enzymes are required. We need to develop a different method for measuring the energy of this reaction, and this requires more experimentation.

To begin our observations, we will work with a few reactions for which we can measure the energy change. Hydrogen gas, which is of potential interest nationally as a clean fuel, can be generated by the reaction of carbon (coal) and water:

Calorimetry reveals that this reaction requires the input of 90.1 kJ of heat for every mole of C(s) consumed.

It is interesting to ask where this input energy goes when the reaction occurs. One way to answer this question is to consider the fact that Reaction (3) converts one fuel, C(s), into another, H ₂ (g). To compare the energy available in each fuel, we can measure the heat evolved in the combustion of each fuel with one mole of oxygen gas. We observe that