Chemical Thermodynamics: Enthalpy - Online Tutor, Practice Problems & Exam Prep

Enthalpy, denoted as ΔH, measures heat transfer in a system at constant pressure. Exothermic processes release heat, resulting in a negative ΔH, while endothermic processes absorb heat, leading to a positive ΔH. For example, condensation and freezing are exothermic, while melting and vaporization are endothermic. Energy diagrams illustrate these changes, showing reactants and products at different energy levels. The change in enthalpy can be calculated using the equation: $\Delta H=\mathrm{Products}-\mathrm{Reactants}$.

Enthalpy

Hey everyone, so enthalpy represents the amount of kinetic energy in the form of heat between a system and its surroundings under constant pressure. And we're going to say here that our standard enthalpy of formation, which is ΔH°f, represents the heat change between reactants and products that are in their standard or natural states. Now, with this whole idea of enthalpy, we have 2 other concepts, that of exothermic and endothermic processes. Now if we took a look at exothermic processes, we could take a look at an exothermic solution. Here we would have a sign for q, which is heat. That is negative, and since we're under constant pressure, enthalpy which is ΔH would also have the same sign, so it would also be negative. Now when it comes to exothermic processes, here we have 3 moles of hydrogen gas reacting with 1 mole of nitrogen gas to produce 2 moles of ammonia gas. And when we're talking about exothermic processes, we're going to say that we start off initially with weak bonds for our reactants, and they help to make strong bonds with our product. Now, here we're talking about weak lattice bond energy, so it's easier to break those bonds apart so they can reassemble and form stronger product bonds later on. Now, why does it do this? Well, we're going to say when it comes to an exothermic process, we're going to release heat in order to form bonds. So if you think about it, let's say you have gas molecules inside of a container and they're bouncing everywhere within the container. They're moving so fast and are so far from each other that they're not able to connect. But what happens if the gas molecules start to lose their heat? They begin to slow down, slow down enough that they will condense into a liquid and thereby be closer. Keep driving heat away from these molecules and it'll eventually solidify, thereby forming strong connections with one another. So here if our gas molecules are releasing heat, that heat has to go somewhere. If we were to touch the container in which the gas molecules reside, the container would feel warm to the touch. That's because the gas molecules are releasing heat and my hands, the container, are absorbing that heat or feeling that heat. Now here, if we're talking about an exothermic process, we have our phase changes. We can go from gas to liquid, which would be condensation. We can go from liquid to solid which is freezing, or we can go straight from gas to solid which is deposition. Now, if we look at the opposite process, we have an endothermic process. Since it's an opposite process, the signs of q and H here would be positive. Here we have HBr breaking into 2 ions, H+ ion and Br− ion. With an endothermic process, we're going to say we initially have strong lattice bond energy. So the connection between the molecules or atoms at the beginning is pretty strong and requires some outside energy source to break those bonds. As a result of this, we're going to have weak bonds formed for the products. Endothermic processes, they absorb heat, and they absorb heat to break bonds. Think of it like an ice cube. You have an ice cube which is solid water. Let's say that I were to heat up that ice cube. The ice cube molecules would absorb this external heat and use it to break themselves free, thereby transitioning from a solid to a liquid. If I kept applying additional heat, there could come a point where the liquid eventually starts to vaporize into a gas. Now, so I am plugging in heat into my solid molecule here. It's taking in heat from the outside. So if I were to touch the container, the container would feel cold to the touch. That's because the molecules are absorbing the external energy from the solution around it, from my hand, from wherever, in order to gain enough energy to break its bonds and free itself. So if we're looking at the phase changes associated with this, we can say that we're going from a solid to a liquid that's melting or another term for it is fusion. If we're going from a liquid to a gas then that's vaporization, or we can go straight from a solid to a gas which is sublimation. Now in terms of energy diagrams, the layout for thermochemical equations can be displayed by the images below. For the first one, we have enthalpy and we're going to say here that we start off with our reactants and we can see that at the end, our products are at a lower state. So we're going to say they're energetically at a lower state. To do this, the reactors would have to release their excess energy. So since we're releasing heat, this will be negative for ΔH which will mean it's exothermic. In the next image, we have reactants down here and then products are higher up here. How did this happen? The reactants had to absorb energy, and because they're absorbing energy, we're going to create products that are at a higher energy state. So, we're absorbing, so ΔH here is positive and therefore it's endothermic. And then finally, we tend not to talk about this term but it is involved as well. It could be a point where our reactants and products are at the same levels. Yes. There was an increase in energy at some point in the reaction, but at the end, reactants and products are at the same point. Here the change in our enthalpy will be 0 and this would be called a thermo neutral reaction. Now talking about ΔH being positive or negative or 0 really hinges on the fact that when we're looking at energy diagrams, this change in our enthalpy of formation can be attributed to this equation of products minus reactants. Here, I'm showing it in a fancier way that it is the summation of the enthalpies of formation of all the products minus the summation of the enthalpies of formation of all our reactants. But in layman's terms, it's just products minus reactants. So just remember with enthalpy we're talking about the energy involved in the transferring of heat, between our systems and our surroundings under constant pressure. Here, this can result in an endothermic process, an exothermic process, or even a thermal neutral process.

Here’s what students ask on this topic:

What is the difference between exothermic and endothermic processes in terms of enthalpy?

Exothermic processes release heat to the surroundings, resulting in a negative change in enthalpy (ΔH < 0). This means the products are at a lower energy state than the reactants. Examples include condensation and freezing. Endothermic processes, on the other hand, absorb heat from the surroundings, leading to a positive change in enthalpy (ΔH > 0). This indicates that the products are at a higher energy state than the reactants. Examples include melting and vaporization. The sign of ΔH helps determine whether a process is exothermic or endothermic.

How do you calculate the change in enthalpy (ΔH) for a chemical reaction?

The change in enthalpy (ΔH) for a chemical reaction can be calculated using the equation:

$\Delta H=\frac{\sum H\left(\mathrm{products}\right)}{-\sum H\left(\mathrm{reactants}\right)}$

This equation means you sum the enthalpies of formation of all the products and subtract the sum of the enthalpies of formation of all the reactants. The result will give you the overall change in enthalpy for the reaction.

What is standard enthalpy of formation (ΔH_f°) and how is it used?

The standard enthalpy of formation (ΔH_f°) is the change in enthalpy when one mole of a compound is formed from its elements in their standard states under standard conditions (298 K and 1 atm). It is used to calculate the enthalpy changes in chemical reactions by summing the standard enthalpies of formation of the products and subtracting the sum of the standard enthalpies of formation of the reactants. This helps in determining whether a reaction is exothermic or endothermic.

What are the phase changes associated with exothermic and endothermic processes?

Exothermic phase changes release heat and include:

• Condensation (gas to liquid)
• Freezing (liquid to solid)
• Deposition (gas to solid)

Endothermic phase changes absorb heat and include:

• Melting (solid to liquid), also known as fusion
• Vaporization (liquid to gas)
• Sublimation (solid to gas)

These phase changes illustrate how heat transfer affects the state of matter.

How do energy diagrams illustrate exothermic and endothermic reactions?

Energy diagrams show the energy levels of reactants and products. In an exothermic reaction, the reactants start at a higher energy level, and the products are at a lower energy level, indicating that energy is released (ΔH < 0). In an endothermic reaction, the reactants start at a lower energy level, and the products are at a higher energy level, indicating that energy is absorbed (ΔH > 0). These diagrams help visualize the energy changes during a reaction.