1.1 - Enthalpy Change and "Heat" (Q)
Module 4
Lesson 1.1 Enthalpy Change and "Heat" (Q)
Key Concepts
In order to measure enthalpy change, we must infer changes in a chemical system's potential energy. But how do we determine whether a change in potential energy has occurred?
If potential energy is lost by a chemical system (as the bonds rearrange), this loss is equivalent to the kinetic energy gained by the surroundings. Because of this energy gain, the temperature of the surroundings will increase.
However, if potential energy is gained by a chemical system, this gain will be equal to the kinetic energy lost by the surroundings. Because of this energy loss, the temperature of the surroundings will decrease.
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In chemistry, heat does not refer to temperature. Temperature measures average kinetic energy. Heat, on the other hand, is the flow of energy from a warmer area to a cooler area. Heat is represented as Q and is measured in either joules (J) or calories. |
In other words, enthalpy change (ΔH) can be inferred from the size of a thermal energy change in the surroundings, as measured by temperature change (of the surroundings). If the surroundings become warmer, it means that the system has lost heat (exothermic). Therefore we can infer that the system has lost potential energy. If, however, the surroundings become cooler, it means that the system has gained heat (endothermic). In this case we can infer that the system has gained potential energy.
ΔH = Q (thermal energy transferred or "heat")
If the chemical reaction is exothermic, ΔH is negative.
If the chemical reaction is endothermic, ΔH is positive.
Watch
The concepts of exothermic and endothermic can be confusing. Distinguishing between "system" and "surroundings" is very important.
View this YouTube video to help clarify!
© Tyler DeWitt