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Why can we measure enthalpy changes, but not absolute enthalpies?
- What would be considered the simplest explanation for high school students regarding why enthalpy cannot be measured, although changes in enthalpy can be?
At this level, just explain that "enthalpy change" is
a synonym for "heat at constant pressure". For example, if an object absorbs 100 J of heat at 1 atm, you can say "the enthalpy of this object increased by 100 J." But you really know nothing about the value of the enthalpy before or after the heat was absorbed.
A metaphor using a change in a more familiar state property that is impractical or impossible to measure absolutely
is very helpful. For example:
- You buy groceries worth $42.00. You can't say what your net financial worth is at that instant, down to the penny; but you can tell they're $42.00 less than before you bought groceries. It doesn't really matter what your net worth is; nothing interesting
happens unless you spend money, or earn money.
- You climb to the top of a building that is 203 feet high. You don't know how far away the center of the earth is, but you know it's 203 feet further away than it was before you climbed the building.
- You dump a liter of water into the ocean. You don't know what the volume of the ocean is, but you know it's a liter more
than it was before you came along.
Note that you're free to define a reference point, though, and measure enthalpies relative to that
reference point (as we do with heats of formation).
At a more advanced level, enthalpy is defined in terms of internal energy: H = U + PV.
The definition was made solely to provide a convenient state function for constant pressure calorimetry (so that H = qP).
Author: Fred Senese firstname.lastname@example.org