 # Endergonic and Exergonic vs Endothermic and Exothermic

One of the most common things I come across when working with students studying for the 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:

Entropy, S

For our discussion, Entropy is a measurement of disorder.

Enthalpy, H

The potential energy that is 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 (transformation) 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 under constant pressure conditions circa the first law of thermodynamics and the definition of enthalpy.

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. Note how this relationship is specifically expansion or compression against a constant external pressure. image credit to University of Texas

Lets combine the first law of thermodynamics with the definition of enthalpy: Enthalpy changes can be associated with the flow of heat under constant pressure circumstances such as inside of our cells.

So the flow of heat (enthalpy) is the flow of potential energy into and out of the system. in exo/endothermic reactions POTENTIAL energy is flowing out and in.

Measuring potential energy changes is great and all, but how much of this potential energy can we actually use to do something with ? The energy you can use to do work is basically the potential energy released from the reaction subtracted by the energy lost via entropy. Temperature is mostly just there to make everything dimensionally consistent.

The 2nd law of thermodynamics tells us that we cant use ALL of the energy in a chemical reaction to do work (some will be lost due to entropy), only a small amount of that potential energy can be used to do things such as move a piston. So chemists had to come up with Endergonic and Exergonic to explain changes in usable (free) energy. Note that energy can be absorbed or released but what changes is the amount of GIBBS FREE energy ie energy you can use to do work in endergonic and exergonic reactions.

TLDR:

-Exo/Endotehrmic we are measuring changes in potential energy states

-cant use all potential energy to get work done (entropy causes us to lose energy)

-gotta measure energy we can use for work as energonic and exergonic