5.4 Thermochemistry: calorimetry  (Page 5/14)

This link shows the precipitation reaction that occurs when the disk in a chemical hand warmer is flexed.

Heat flow in an instant ice pack

When solid ammonium nitrate dissolves in water, the solution becomes cold. This is the basis for an “instant ice pack” ( [link] ). When 3.21 g of solid NH ₄ NO ₃ dissolves in 50.0 g of water at 24.9 °C in a calorimeter, the temperature decreases to 20.3 °C.

Calculate the value of q for this reaction and explain the meaning of its arithmetic sign. State any assumptions that you made.

Solution

We assume that the calorimeter prevents heat transfer between the solution and its external environment (including the calorimeter itself), in which case:

$q_{\text{rxn}}=\text{−}{q}_{\text{soln}}$

with “rxn” and “soln” used as shorthand for “reaction” and “solution,” respectively.

Assuming also that the specific heat of the solution is the same as that for water, we have:

$\begin{array}{}\\ q_{\text{rxn}}=\text{−}{q}_{\text{soln}}={-(c\times m\times \text{Δ}T)}_{\text{soln}}\\ =\text{−[(}4.184\text{J/g °C})\times (53.2\text{g})\times (20.3\text{°C}-24.9\text{°C})]\\ =\text{−[(}4.184\text{J/g °C})\times (53.2\text{g})\times (\mathrm{-4.6}\text{°C})]\\ \text{+}1.0\times {10}^3\text{J}=\mathrm{+1.0}\text{kJ}\end{array}$

The positive sign for q indicates that the dissolution is an endothermic process.

Check your learning

When a 3.00-g sample of KCl was added to 3.00 $\times$ 10 ² g of water in a coffee cup calorimeter, the temperature decreased by 1.05 °C. How much heat is involved in the dissolution of the KCl? What assumptions did you make?

Answer:

1.33 kJ; assume that the calorimeter prevents heat transfer between the solution and its external environment (including the calorimeter itself) and that the specific heat of the solution is the same as that for water

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If the amount of heat absorbed by a calorimeter is too large to neglect or if we require more accurate results, then we must take into account the heat absorbed both by the solution and by the calorimeter.

The calorimeters described are designed to operate at constant (atmospheric) pressure and are convenient to measure heat flow accompanying processes that occur in solution. A different type of calorimeter that operates at constant volume, colloquially known as a bomb calorimeter    , is used to measure the energy produced by reactions that yield large amounts of heat and gaseous products, such as combustion reactions. (The term “bomb” comes from the observation that these reactions can be vigorous enough to resemble explosions that would damage other calorimeters.) This type of calorimeter consists of a robust steel container (the “bomb”) that contains the reactants and is itself submerged in water ( [link] ). The sample is placed in the bomb, which is then filled with oxygen at high pressure. A small electrical spark is used to ignite the sample. The energy produced by the reaction is trapped in the steel bomb and the surrounding water. The temperature increase is measured and, along with the known heat capacity of the calorimeter, is used to calculate the energy produced by the reaction. Bomb calorimeters require calibration to determine the heat capacity of the calorimeter and ensure accurate results. The calibration is accomplished using a reaction with a known q , such as a measured quantity of benzoic acid ignited by a spark from a nickel fuse wire that is weighed before and after the reaction. The temperature change produced by the known reaction is used to determine the heat capacity of the calorimeter. The calibration is generally performed each time before the calorimeter is used to gather research data.