So far in this unit, we have learned that energy is required to break bonds and that energy is released when bonds are made, but how much energy? How can we actually measure the energy transferred in one or both of these processes?

Because the energy transfer involved in both breaking and making bonds involves heat, or thermal energy, we can employ a method called calorimetry to measure the heat transfer. Calorimetry is the science of measuring heat. In aqueous solutions, the energy is transferred to or taken away from the water, therefore, the thermal energy transferred may be calculated using the change in temperature of the water itself.

A calorimeter can be used to measure the amount of thermal energy transferred when bonds are broken or formed. A calorimeter will consist of a container with a cover, and a thermometer or temperature probe. The ideal calorimeter (for our purposes) must prevent the transfer of thermal energy between the calorimeter's container and its environment. This allows for an accurate determination of the thermal energy that has either been transferred to or from the water within the calorimeter.

Experiment with the simulation by observing two different chemical processes.

• The separation of the ionic compound potassium iodide (KI) in water:

KI (s) → K⁺ (aq) + I⁻ (aq)

• The synthesis of calcium hydroxide, Ca(OH)₂ , and subsequent production of hydrogen gas when solid calcium metal (Ca) is added to water:

Ca(s) + 2 H₂O (l)→ Ca(OH)₂ (aq) + H₂ (g)

Begin by hitting "setup” and then “run/pause”. In order to monitor the temperature of your experiment you will need a sensor in the container. Press “place sensor” in order to add a sensor inside your container with your cursor.

In order to observe the ionic compound separate in water, you will need to press the "Add KI" button; to observe the synthesis of calcium hydroxide and production of hydrogen gas, you will need to press the "Add Ca" button.

Within this unit, we have observed changes in thermal energy by observing changes in temperature within a closed system. We saw this in a temperature decrease with ionic solids dissolving in water in lessons 2 and 3. We also saw this with a temperature increase through the synthesis of an ionic compound in lesson 4.

The thermal energy transferred when bonds are broken or formed is referred to as the enthalpy change (∆H). For our purposes, energy transfer will always occur within a closed system.

Because any energy transfer discussed in this class will occur within a closed system, the following relationship will be true for these energy transfers:

Enthalpy change = heat energy

∆H = q

In other words, as long as the energy is transferred within a closed system the enthalpy change (∆H) will be equal to the heat energy (q) that is transferred:

q = m c ∆T

∆H = m c ∆T

where:

q = heat energy in joules (J), so ∆H = energy transferred in joules (J)

m = mass of sample in grams (g)

c = specific heat (J/g^OC)

ΔT = change in temperature (final temperature - initial temperature)

Specific heat, (c), is an intensive physical property that has a different, characteristic value for every substance. For example, the specific heat for water is 4.18 J/g^OC. This would mean that it will take 4.18 joules of energy to raise the temperature of 1 gram of water 1^OC.

Change in enthalpy is always relative to the energy within a system. If the system gains energy, change in enthalpy (∆H) will be positive. If the system loses energy, the change in enthalpy (∆H) will be negative.

In lessons 2 and 3, we observed the dissolving of an ionic compound in water. A decrease in temperature of the surroundings was observed, and this was labeled as endothermic. In an endothermic reaction, the change in enthalpy (∆H) equals the energy in joules absorbed by the system from the surroundings. Because energy is absorbed into the system, ∆H is positive.

In lesson 4, we observed the synthesis of an ionic compound. An increase in temperature of the surrounding was observed, and this reaction was labeled as exothermic. In an exothermic reaction, the change in enthalpy (∆H) equals the energy in joules released by the system to the surroundings. Because energy is released into the surroundings, ∆H is negative.

The simulation you were working with on page 2 models a calorimeter. The simulation allows you to choose the material to use for your calorimeter using the dropdown menu labeled "container-material".

Begin by hitting "setup” and then “run/pause”. Adjust the "sensor-radius" slider bar to 48, then press “place sensor” to add this to the inside of the container ensuring the sensor does not contact to walls of the container. Now adjust the "sensor-radius" slider bar to 4, then press “place sensor” to add this to the outside of the container, again ensuring the sensor does not contact to walls of the container.

You will choose one of the chemical processes available (either the separation of the ionic compound potassium iodide (KI) in water or the synthesis of calcium hydroxide, Ca(OH)₂ , and subsequent production of hydrogen gas when solid calcium metal (Ca) is added to water) to determine the ideal material to use for your calorimeter. In order to observe the ionic compound separate in water, you will need to press the "Add KI" button; to observe the synthesis of calcium hydroxide and production of hydrogen gas, you will need to press the "Add Ca" button.

Now that you have experimented using several different container materials, you will determine which material is best suited for a calorimeter.

On page 1 of this lesson, it was stated that the ideal calorimeter must prevent the transfer of thermal energy between the calorimeter's container and its environment in order to allow for an accurate determination of the thermal energy transferred.

The data you acquired for aluminum, glass, and Styrofoam is shown below.