Enthalpy — Definition & Overview

When I was in college, I studied environmental chemistry. If you love chemistry and want to be outside, it's the way to go. But, I also liked it because it simplified thermodynamics. In the environment, everything is under constant pressure. That's atmospheric pressure. But, even in the lab, we do many reactions in an open beaker! Any chemical reaction that is open to the air has a constant pressure! When pressure is constant, thermodynamics is more manageable.

Energy Change at Constant Pressure: Enthalpy

We have a particular term for the energy change at constant pressure. It's called enthalpy! We studied internal energy. It was the total energy within a system. That's the sum of the energies from molecular movements, bonds, and intermolecular forces. But what happens when we have an open system? Our system has to push against the atmosphere. So, its volume can change. Enthalpy accounts for the energy of pushing against the atmosphere. So, enthalpy is the internal energy plus the energy that pushes against surroundings'. We define it as $H= U + PV$. The change in enthalpy is $\Delta H= \Delta U +P\Delta V$.

We can simplify the enthalpy equation.

$H= U + PV \rightarrow H= Q + W + PV \rightarrow H= Q - PV +PV \rightarrow H=Q$.

Most often, we define enthalpy at a constant pressure. So, we write the equation $H=Q_{p}$.

There are a few cases where we have constant volume. For example, if all the reactants and products are solids or liquids. They don't compress at atmospheric pressure. So, their volumes don't change. These cases are rare. But, we can denote constant volume enthalpy as $H=Q_{v}$.

What Is PV-Work?

Imagine you have a chemical reaction. It's open to the air. One of the products is a gas. How does the gas escape the beaker? It does work against the atmosphere! For the gas to diffuse, it has to collide with atmospheric gases. The collisions exert a force. So, the gas is doing work. Any expansion or compression caused by a gas is pressure-volume work.

It might be easier to imagine a moveable piston. Remember, a movable piston keeps the pressure constant. As the gas fills the reaction chamber, it pushes the piston. The gas applied a force to the piston. It did work.

What's an example of PV-work? A car engine! Inside a car engine, a combustion reaction converts gasoline to carbon dioxide. The gas pushes against the cylinder. As the cylinder moves, it rotates the crankshaft. And your car moves forward.

Why Do We Need Enthalpy?

We've seen that at constant pressure, enthalpy equals heat. Why does this help us? The problem is that heat is a path function. Remember, heat is the energy transferred because of temperature differences. Heat transfer is a mechanism. It depends on how it's done. Let's think about it mathematically. We know $U= Q + W$. Let's say systems' internal energy was 10.0 Joules. We could have many combinations of Q and W. For example, Q=10.0 J and W= 0.0 J or Q= 5.4J and W=4.6J. But, the internal energy is a state function. In our example, it is always 10.0 J. So, different paths of Q and W led to the same internal energy.

So, what was the benefit of defining enthalpy as $\Delta H= \Delta U +P\Delta V$? U, P, and V are all state functions. So, when we simplified to $\Delta H= Q_{p}$, we created a state function. When our pressure is constant, we can calculate enthalpy for any state. We'll see that enthalpy relates to temperature, the material, and the material's mass. When we know these, we can find the internal energy's Q-term.

(Video) Enthalpy

by Ben's Chem Videos

Enthalpy is a positive or negative change in heat released by a system under constant pressure. In this video, Ben describes this idea more in detail and shows you how to derive the equation for enthalpy.