THC_Calculating Energy Change - Chemical Reactions Lesson
THC_Calculating Energy Change - Chemical Reactions
As you learned on the previous page, the energy change associated with a chemical reaction is expressed by heat of reaction, ∆H. These values can be can be calculated a couple of ways.
Hess' Law
Remember that when the direction of a thermochemical equation is reversed, the sign of the enthalpy of the reaction is reversed. When an equation is doubled, tripled, cut in half, etc., the quantity of ∆H changes by that same ratio. When reactions are mathematically manipulated in this way, we are actually using the concept known as Hess' Law. Hess' Law says that the total enthalpy of a reaction is independent of the reaction pathway. This means that if a reaction is carried out in a series of steps, the enthalpy change (∆H) for the overall reaction will be equal to the sum of the enthalpy changes for the individual steps.
To better understand this concept, think about climbing a mountain. To get to the top, you can take the road that goes around the mountain. Or, you can hike straight up the side.
Regardless of the path you choose, your change in elevation will be the same. ∆H is the same way in that its value is independent of the path taken to establish its value.
The advantage of Hess Law is that it allows us to calculate the enthalpy change for a chemical reaction in which the enthalpy data is not directly known.
Watch the following video to see how to use Hess' Law to calculate ∆Hrxn.
Hess' Law can be shown visually through an enthalpy diagram. The one shown here represents the above reaction. The change in enthalpy, ΔH, is represented by the distance between the horizontal lines. It is important to note that we cannot know the actual H values (the horizontal lines) for any chemical since it is impossible for us to measure total heat content, H. But, through calorimetry, we can measure ∆Hrxn.
Standard Heat of Formation
Another useful tool to calculate heats of reaction is called standard enthalpy of formation or standard heat of formation, ΔHf °. ΔHf ° is a measure of the energy released or consumed when one mole of a substance is created under standard conditions (1 atm and 25oC) from its pure elements.
Look at the thermochemical equations below.
Only the first reaction represents a standard heat of formation reaction. This is because only the first equation produces 1 mole of product from its elements in their standard states (meaning the phase of matter that they are in at 1 atm and 25oC). The third reaction produces 1 mole, but the reactants are not in their standard states since they are not written as the diatomic elements that they are. The second reaction produces 2 moles of product, and therefore is not a standard heat of formation reaction. Note that we often have to use fractions to balance these reactions in order to produce only 1 mole of the product.
You will find a complete list of standard heats of formation in the appendix of your book. Go ahead and look at this chart now and notice the following:
- ΔHf ° values have units of kJ/mol because each is the enthalpy change for the formation of one mole of a chemical species.
- The ΔHf ° value for an element in its standard state is 0 kJ/mol.
- Most ΔHf ° values are negative, meaning that for most species, the formation from elements in their standard states is an exothermic process.
So, now that you understand what heats of formation are, why are they so important? The answer is that we can use them as a shortcut to solve for ∆H of a specific reaction. Yes, we can use Hess' Law and manipulate a variety of reactions to come up with the new reaction. But, knowing the standard heats of formation for many chemicals allows us this shortcut to Hess' Law.
This equation might not look like a shortcut in this form. Just think of it as products minus reactants. Look at this example problem to see how to use this formula.
Remember to work on the module practice problems as you complete each section of content.
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