One of the keys to thermodynamics is understanding state functions. The practical application is seeing that if you know information about the initial and final state of a system, you can determine the change. Sometimes, it is too complicated to directly measure that change, but we can measure smaller steps in the process and combine them to determine the overall change. This is Hess’s Law.
- Use Hess’s law to determine enthalpy changes for reactions. (5.6)
- Use standard enthalpies of formation to calculate ΔH° for reactions. (5.7)
In addition to the class discussion, you should familiarize yourself with the key ideas of this topic. These slides relate to sections 5.6-5.7 of your textbook.
Now that you are familiar with these concepts, you should put them into practice.
When you have finished working out the practice problems, take a picture of your work and add it to your OneNote.
Complete the problems below, then take a picture of your work and post it in OneNote. Be prepared to present these problems in class.
Given the data
|ΔH = +180.7 kJ/mol|
|ΔH = -113.1 kJ/mol|
|ΔH = -163.2 kJ/mol|
Use Hess's law to calculate ΔH for the reaction
Many cigarette lighters contain liquid butane, . Using standard enthalpies of formation, calculate the quantity of heat produced when 5.00 g of butane is completely combusted in air under standard conditions.
The precipitation reaction between AgNO3(aq) and NaCl(aq) proceeds as follows:
AgNO3(aq) + NaCl(aq) → NaNO3(aq) + AgCl(s)
(a) By using Appendix C, calculate ΔH° for the net ionic equation of this reaction.
(b) What would you expect for the value of ΔH° of the overall molecular equation compared to that for the net ionic equation? Explain.
(c) Use the results from (a) and (b) along with data in Appendix C to determine the value of ΔHׄ°f for AgNO3(aq).
2002 FRQ #5
Answers to these mastery problems can be found in the Content Library within OneNote.
Three for three!