Textbook 4D.3

[BKoh_2E]

The reaction of 1.40 g of carbon monoxide with excess water vapor to produce carbon dioxide and hydrogen gases in a
bomb calorimeter causes the temperature of the calorimeter assembly to rise from 22.113 8C to 22.799 8C. The calorimeter
assembly is known to have a total heat capacity (calorimeter constant) of 3.00 kJ/C?. (a) Write a balanced equation for the
reaction. (b) Calculate the internal energy change, Delta U, for the reaction of 1.00 mol CO(g).

In the solution manual, the answer for part b is -41.2kj/mol. I thought the solution would be positive 41.2kJ/mol because the temperature rose from 22.113C to 22.799C, which means heat was absorbed and that would mean it was an endothermic reaction. Or do calorimeters only rise in temperature when it measures heat released from exothermic reactions?

[Keerthana Sundar 1K]

I believe it's because the calorimeter is the surroundings and the reaction itself is the system. If the calorimeter warms up, then that means the reaction has released heat and the calorimeter has absorbed that heat released from the reaction. If the calorimeter cools down, then that means the reaction has absorbed heat from the calorimeter. In the case of this problem, we are viewing the enthalpy change regarding the reaction itself, not the calorimeter. Let me know if I'm wrong or if anyone has anything else to add. Hopefully that helped!

[Kaitlin Eblen 1I]

I completely agree with Keerthana. Here is a diagram I found a helpful diagram that helps clarify how a calorimeter works. So an increase in temp. of the solution (recorded by the calorimeter) would

I completely agree with Keerthana. Here is a diagram I found that helps clarify how a calorimeter works. So an increase in temp. of the solution (recorded by the calorimeter) would mean that an exothermic reaction occurred (negative delta q).

[Shivani Sakthi 1l]

Hi!
I completely agree with the responses above. It is important to think of this model of the calorimeter and the carbon monoxide reaction as a system and its surroundings. The reaction mixture is the SYSTEM while the calorimeter is the SURROUNDINGS. We are sure that the calorimeter is considered the surroundings because it is a location in which the reaction is immersed. We are told that the temperature of the calorimeter, or the SURROUNDINGS rises. Therefore, we know that heat is being transferred from the reaction to the calorimeter. This means that the reaction is exothermic, and therefore, the internal energy change of the reaction would be negative. Since the problem states "calculate the internal energy change, Delta U, for the reaction of 1.00 mol CO(g)", we know that we are looking at the enthalpy for the reaction of 1.00 mol of CO, not in terms of the calorimeter.

This is very similar to the example Professor Lavelle Discussed in lecture. For example, if our system is the reaction of OH- and H+ when mixing an acid and base, heat is released from the reaction to the SURROUNDINGS, which are water molecules. So, if the water molecules increase in temperature, we know that the reaction is giving off heat into the surroundings. This means the reaction is exothermic. However, the heat that the system gives off is equal and opposite to its surroundings. So, the delta H of the surrounding water molecules would be endothermic, since it is absorbing energy, but the same value of the reaction. this same principle can be applied to the context of the CO reaction and the calorimeter.