Heat, Energy and now Enthalpy? What is going on? Our episode starts by defining and distinguishing between these concepts.
Heat, Energy and now Enthalpy? What is going on? Our episode starts by defining and distinguishing between these concepts (0:54) before taking a closer look at enthalpy itself. The episode connects the mathematical sign of enthalpy to the heat being absorbed / released (1:39) and discusses the extensive character of enthalpy (2:06) as well the value of reversed reactions (2:23) and its dependency on the state (2:47). We also recap two approaches to calculate the enthalpy of a reaction: taking into account the bonds broken and bonds formed (3:06) as well as the standard enthalpies of formation (5:25). To clarify the calculations using the standard enthalpies of formation, we discuss the combustion of propane (7:10).
Question (9:20): What is the law called that allows us to use the sum of enthalpy changes, independent of the steps?
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Hi and welcome to the APsolute Recap: Chemistry Edition. Today’s episode will recap enthalpy.
Lets Zoom Out:
Unit 6 - Thermodynamics
Topics 6.6 - 6.8 - Enthalpy
Big idea - Energy
So far we have recapped internal energy as well as heat. Now a third player enters the arena: Enthalpy! What is going on?! Are these all the same? Where is the difference? Deep breath, we’ve got you!
Let’s zoom in:
Whenever we talk about Energy, we are referring to “internal energy”, which is defined as the sum of all kinetic and potential energies of the components of the system. Energy can be gained or released in the form of heat or work done on or by the system. In chemistry, the work is Pressure x Volume. Enthalpy is defined as the internal energy plus the product of pressure and volume. Most of the reactions we are looking at in AP Chemistry happen under constant pressure. Therefore, substituting heat and work into the equation, Enthalpy is proportional to q, heat. They are different concepts, but, in AP Chemistry, you do not have to distinguish between internal energy and enthalpy.
Phew, now that we got this out of the way, let’s take a closer look. When we discuss the energetics of a chemical reaction in the form of heat, we look at the change of enthalpy between products and reactants. When heat is being released from the system to the surroundings, the value of enthalpy is negative. When heat is being absorbed by the system, the value of enthalpy is positive. That already sounds familiar - we’ve discussed the terms exothermic and endothermic before!
Enthalpy is an extensive property. Reminder: extensive properties are properties that depend on the amount of matter. That means, the amount of heat released with the combustion of 1 mole of methane doubles when I combust 2 moles of methane gas. If I reverse a reaction, then the enthalpy remains equal in size, but with the opposite mathematical sign. For example, if the forward reaction releases 750 kJ and therefore is written as negative 750 kJ, then the reverse reaction absorbs 750 kJ and is written as positive 750 kJ. Enthalpy also depends on the state of the reactants and products. If my reaction produces liquid water it will release more energy than if it would produce water vapor because the conversion from liquid water to water vapor takes energy and therefore reduces the amount of heat that is being released.
To calculate the change of enthalpy of a chemical reaction and therefore the amount of heat being absorbed or released, we can look at the energy content of the products and the reactants - the difference between those two will be our change in enthalpy. But what happens “in between”? As we’ve discussed in our episode about chemical and physical changes, during chemical reactions bonds are broken, which requires energy, and new bonds are being formed, which releases energy. So, to determine the change in enthalpy, we can use the average bond energies. Bond energies are tabulated values that show us how much energy has to be absorbed to break a bond and how much energy is being released when a new bond forms. As discussed in an earlier episode, this depends on different factors, such as bond order and bond length.
To determine the change in enthalpy, we add up the energy it takes to break all the bonds as well as the energy it releases when the new bonds are formed and calculate “bonds broken minus bonds formed”. If the energy released is greater than the energy absorbed, our change in enthalpy has a negative value and the reaction is exothermic. If the energy absorbed is greater than the energy released, the change in enthalpy is a positive value, indicating an endothermic reaction.
Another way of calculating the change in enthalpy is using the standard enthalpies of formation, which are tabulated values. These are defined as the enthalpy change for the reaction in which one mole of a compound is made from its constituent elements in their elemental forms. What does that mean? The standard enthalpy of formation tells us that if we are forming 1 mole of CO2 from its elements, solid carbon and gaseous oxygen, at 1 atm and 25 degrees Celsius the reaction releases 393.5 kJ. The 1 atm and 25 degrees are the “standard” conditions. In the calculation itself, the standard enthalpies of elements in their standard states is zero. In our example, the standard state of oxygen is the diatomic oxygen, not ozone and the standard state of carbon is graphite and not diamond.
How does that help us when it comes to the change of enthalpies of a reaction? We can use these values to calculate ANY reaction by breaking down our reaction into smaller steps. In a first step, we could use the reverse of the enthalpies of formation to convert our reactants into their elements. And in the second step we use the standard enthalpies of formation to form our products from the constituent elements. Since we, in most cases, have coefficients, we need to multiply our standard enthalpies of formations by the coefficients. If we now take the difference between the sum of the standard enthalpies of formation of the products minus the sum of the standard enthalpies of formation of the reactants we get our overall change in enthalpy.
Still confused? Let’s make an example: the combustion of 1 mole of propane gas with 5 moles of O2 to form 3 moles of CO2 and 4 moles of H2O. We have two steps: In step one, we decompose 1 mole of propane into its elements, 3 moles of C and 4 moles of diatomic hydrogen. This is the reverse of the formation of propane and therefore will require 103.85 kJ. Oxygen is already an element, so nothing we have to do there.
Now we are forming our products: first, the formation of 3 moles of CO2 from 3 mol of C and 3 moles of diatomic oxygen. According to the standard enthalpies of formation, 1 mole will release 393.5 kJ, so three moles release 1180.5 kJ. The formation of 4 moles of water from its element releases 4 x 285.8 kJ. To calculate the change in enthalpy, we calculate the sum of the products minus the sum of the reactants: -1180.5 + (-1143.2) - (-103.85 kJ) equals -2219.9 kJ. As combustions usually are, this is a very exothermic reaction.
To recap:
In AP Chemistry enthalpy is equivalent to heat. When heat is being released from the system to the surroundings, the value of enthalpy is negative. When heat is being absorbed by the system, the value of enthalpy is positive. Enthalpy is an extensive property and depends on the state of matter. To calculate the change in enthalpy, we can use the bond energies and take the difference between bonds broken and bonds formed. Another approach is using the standard enthalpies of formation and calculating the sum of the standard enthalpies of formation of the products minus the sum of the standard enthalpies of formation of the reactants.
Coming up next on the APsolute RecAP Chemistry Edition: Equilibrium Constant
Today’s Question of the day is about calculating enthalpy.
What is the law called that allows us to use the sum of enthalpy changes, independent of the steps?