One of the most common things I come across when working with students studying for the MCAT, PCAT, DAT, AP, and college level chemistry courses is understanding the differences between endergonic and exergonic vs. endothermic and exothermic.
Before we get into the differences lets point out what they have in common: energy is going somewhere. In both thermic and ergonic processes energy is either going in or out. The major difference being what kind of energy is moving in and out, and the details of that energy movement.
I am going to assume that if you are asking the difference between exo(endo)thermic and exer(ender)gonic you are probably familiarized with the differences between enthalpy, entropy, and gibbs free energy. In case you need a refresher here it is:
For our discussion, Entropy is a measurement of disorder.
The changes in potential energy involved in any process of transformation, such as breaking and forming chemical bonds in a reaction.
For our discussion we can just conflate this with the flow of heat under constant pressure conditions: see below.
Gibbs free energy, G
The energy of a chemical reaction you can use to do work.
When we measure changes such as exothermic or endothermic processes: we are measuring changes in potential energy involved in the formation and breaking of chemical bonds in a particular reaction (exo and endothermic).
AN: If you are really into the detailed differences between enthalpy of reaction from enthalpy of formation, click here. For our discussion ill just identify them as the same since they have little relevance to understanding the definition of exo and endothermic.
Changes in enthalpy: either exothermic or endothermic manifest themselves as the flow of heat ( changes in kinetic energy of particles in our reaction) under constant pressure conditions, and can be measured as changes in potential energy circa the first law of thermodynamics and the definition of enthalpy. ******
So in exothermic and endothermic processes POTENTIAL ENERGY is changing as energy flows in and out.
Measuring potential energy changes is great and all, but how much of the energy loss/gain can we actually use to do something with ?
The 2nd law of thermodynamics tells us that we cant use ALL of the energy in a chemical reaction to do work, only a small amount of it. So chemists had to come up with Endergonic and Exergonic to explain changes in Gibbs Free Energy (the energy you can use to do work) with a chemical reaction.
-Exo/Endotehrmic we are measuring changes in potential energy states
-cant use all potential energy to get work done
-gotta measure energy we can use for work as energonic and exergonic
Why is it okay to equate enthalpy changes with the flow of heat?
The first law tells us that the change in energy for the system is equal to the flow of heat (positive or negative) + the work
This is a very useful metric for predicting what compounds will form under certain conditions and the TOTAL potential energy changes associated with a chemical reaction.
Normally work is equal to force multiplied by distance, we need to do some arithmetic to convert a force and distance relationship into pressure and volume.
Lets combine the first law of thermodynamics with the definition of enthalpy:
The flow of heat (enthalpy) is the flow of potential energy into and out of the system.