Calorimetry - measuring energy changes from combustion Energy can be released in chemical reactions as light, sound or electrical energy. The diagram shows a simple calorimetry experiment to measure the heat energy released from burning fuel: Calorimetry method Cold water is measured into a copper calorimeter a small metal can.
The starting temperature of the water is recorded. The water is heated using the flame from the burning fuel. The final temperature of the water is recorded.
Fair testing When comparing different fuels, it is important to carry out a fair test. Noting that the final temperature of both the rebar and water is A g piece of copper is dropped into mL of water at Calculate the initial temperature of the piece of copper.
Assume that all heat transfer occurs between the copper and the water. Assuming that all heat transfer occurs between the copper and the water, calculate the final temperature.
This method can also be used to determine other quantities, such as the specific heat of an unknown metal. The final temperature is Use these data to determine the specific heat of the metal.
Use this result to identify the metal. Noting that since the metal was submerged in boiling water, its initial temperature was Comparing this with values in Table T4 , our experimental specific heat is closest to the value for copper 0. After 5 minutes, both the metal and the water have reached the same temperature: Determine the specific heat and the identity of the metal.
Note: You should find that the specific heat is close to that of two different metals. Explain how you can confidently determine the identity of the metal. This specific heat is close to that of either gold or lead. It would be difficult to determine which metal this was based solely on the numerical values. When we use calorimetry to determine the heat involved in a chemical reaction, the same principles we have been discussing apply.
The amount of heat absorbed by the calorimeter is often small enough that we can neglect it though not for highly accurate measurements, as discussed later , and the calorimeter minimizes energy exchange with the surroundings. Because energy is neither created nor destroyed during a chemical reaction, there is no overall energy change during the reaction.
This means that the amount of heat produced or consumed in the reaction equals the amount of heat absorbed or lost by the solution:. This concept lies at the heart of all calorimetry problems and calculations. When 5. The density of water in this temperature range averages 0. Assume that the calorimeter absorbs a negligible amount of heat and, because of the large volume of water, the specific heat of the solution is the same as the specific heat of pure water.
Given: mass of substance, volume of solvent, and initial and final temperatures. The mass of the solution is. B Because the solution is not very concentrated approximately 0. The heat flow that accompanies dissolution is thus. Where did this heat come from?
It was released by KOH dissolving in water. This experiment tells us that dissolving 5. Because the temperature of the solution increased, the dissolution of KOH in water must be exothermic. A coffee-cup calorimeter contains Commercial solution calorimeters are also available.
Relatively inexpensive calorimeters often consist of two thin-walled cups that are nested in a way that minimizes thermal contact during use, along with an insulated cover, handheld stirrer, and simple thermometer.
More expensive calorimeters used for industry and research typically have a well-insulated, fully enclosed reaction vessel, motorized stirring mechanism, and a more accurate temperature sensor Figure 3. Before we practice calorimetry problems involving chemical reactions, consider a simpler example that illustrates the core idea behind calorimetry.
Suppose we initially have a high-temperature substance, such as a hot piece of metal M , and a low-temperature substance, such as cool water W. If we place the metal in the water, heat will flow from M to W.
The temperature of M will decrease, and the temperature of W will increase, until the two substances have the same temperature—that is, when they reach thermal equilibrium Figure 4.
Under these ideal circumstances, the net heat change is zero:. This relationship can be rearranged to show that the heat gained by substance M is equal to the heat lost by substance W:. The magnitude of the heat change is therefore the same for both substances, and the negative sign merely shows that q substance M and q substance W are opposite in direction of heat flow gain or loss but does not indicate the arithmetic sign of either q value that is determined by whether the matter in question gains or loses heat, per definition.
In the specific situation described, q substance M is a negative value and q substance W is positive, since heat is transferred from M to W. Heat Transfer between Substances at Different Temperatures A g piece of rebar a steel rod used for reinforcing concrete is dropped into mL of water at The final temperature of the water was measured as Calculate the initial temperature of the piece of rebar. Assume the specific heat of steel is approximately the same as that for iron Table 1 in Chapter 5.
Solution The temperature of the water increases from That heat came from the piece of rebar, which initially was at a higher temperature. The density of water is 1. Noting that the final temperature of both the rebar and water is Check Your Learning A g piece of copper is dropped into mL of water at Calculate the initial temperature of the piece of copper.
Assume that all heat transfer occurs between the copper and the water. Assuming that all heat transfer occurs between the copper and the water, calculate the final temperature. This method can also be used to determine other quantities, such as the specific heat of an unknown metal.
The final temperature is Use these data to determine the specific heat of the metal. Use this result to identify the metal. Noting that since the metal was submerged in boiling water, its initial temperature was Comparing this with values in Table 1 in Chapter 5. Check Your Learning A After 5 minutes, both the metal and the water have reached the same temperature: Determine the specific heat and the identity of the metal.
Note: You should find that the specific heat is close to that of two different metals. Explain how you can confidently determine the identity of the metal. This specific heat is close to that of either gold or lead. It would be difficult to determine which metal this was based solely on the numerical values. When we use calorimetry to determine the heat involved in a chemical reaction, the same principles we have been discussing apply.
The amount of heat absorbed by the calorimeter is often small enough that we can neglect it though not for highly accurate measurements, as discussed later , and the calorimeter minimizes energy exchange with the surroundings. Because energy is neither created nor destroyed during a chemical reaction, there is no overall energy change during the reaction.
This means that the amount of heat produced or consumed in the reaction equals the amount of heat absorbed or lost by the solution:. Heat Produced by an Exothermic Reaction When What is the approximate amount of heat produced by this reaction?
Solution To visualize what is going on, imagine that you could combine the two solutions so quickly that no reaction took place while they mixed; then after mixing, the reaction took place. At the instant of mixing, you have The heat given off by the reaction is equal to that taken in by the solution.
It is important to remember that this relationship only holds if the calorimeter does not absorb any heat from the reaction, and there is no heat exchange between the calorimeter and its surroundings. Next, we know that the heat absorbed by the solution depends on its specific heat, mass, and temperature change:. To proceed with this calculation, we need to make a few more reasonable assumptions or approximations. Since the solution is aqueous, we can proceed as if it were water in terms of its specific heat and mass values.
The density of water is approximately 1. The specific heat of water is approximately 4. Substituting these values gives:. The negative sign indicates that the reaction is exothermic. It produces 2. Check Your Learning When mL of 0. How much heat is produced by this precipitation reaction? What assumptions did you make to determine your value? When working or playing outdoors on a cold day, you might use a hand warmer to warm your hands Figure 5.
A common reusable hand warmer contains a supersaturated solution of NaC 2 H 3 O 2 sodium acetate and a metal disc. Bending the disk creates nucleation sites around which the metastable NaC 2 H 3 O 2 quickly crystallizes a later chapter on solutions will investigate saturation and supersaturation in more detail. If the hand warmer is reheated, the NaC 2 H 3 O 2 redissolves and can be reused. Another common hand warmer produces heat when it is ripped open, exposing iron and water in the hand warmer to oxygen in the air.
Salt in the hand warmer catalyzes the reaction, so it produces heat more rapidly; cellulose, vermiculite, and activated carbon help distribute the heat evenly. Other types of hand warmers use lighter fluid a platinum catalyst helps lighter fluid oxidize exothermically , charcoal charcoal oxidizes in a special case , or electrical units that produce heat by passing an electrical current from a battery through resistive wires. 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. When 3. Calculate the value of q for this reaction and explain the meaning of its arithmetic sign.
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