What is the difference between heat and temperature?

use the chapter 8 notes 1.In-class discussion on page 2 and in class discussions on bottom of page 3 to do question 2

you can chose whichever four to do

for the group report question one i would prefer you do problem b

Requirements: you can see in document

Ch. 8: Thermochemistry Aug. 30 – Sep. 8, 2023

Topic 0: Overview and Definitions for heat of reaction: Read through sections 8.1, 8.2, 8.4 & 8.6 and write down definitions for each of the highlighted terms below in your own words.

Section 8.1: What is the difference between heat and temperature? Consider putting a room temperature beaker of water on a hot plate waiting for it to boil. Explain what happens to it in terms of heat and temperature.

Section 8.2: Look up “state function” in this section. Consider heat and temperature from the previous section and state which is a state function and which is not. Consider the beaker of boiling water on a hot plate from the previous section – which is the “system” and which is the “surroundings”? In general, the system will be a chemical reaction of interest and the surroundings will be the aqueous solution and/or the atmosphere surrounding it. If you dissolve NaOH(s) in water, the temperature will increase dramatically. The reaction, NaOH(s) NaOH(aq) is the system and the water (whose temperature we measure) is the surroundings.

Section 8.4: We are interested in the amount of heat given off or absorbed by a chemical reaction. The heat given off or absorbed by the reaction is called enthalpy and denoted . In sections 8.7 – 8.9, we’ll look at different ways of calculating the theoretical enthalpy of reaction. In section 8.5, we’ll look at calorimetry, the technique for measuring (enthalpy change of reaction) experimentally.

Section 8.6: Standard state means that the reaction takes place at 1 atm pressure and is denoted with the degree symbol: . Many reactions are pressure dependent (especially those involving gases), so it’s important to designate that the reaction is carried out at 1 atm pressure. Don’t bother calculating ; it’s silly and not worth your time.

Pre-class questions: What are the definitions of endothermic and exothermic? Which type of reaction will heat the surroundings and which will cool the surroundings?

Homework: 8.36, 8.55, due 8/30/23.

Topic 1, 8/30/23 (Section 8.7): Calculations of heat transfer in Physical and Chemical changes: Let’s do some calculations of enthalpy for physical and chemical changes and interpret what they mean.

Fact: The enthalpy required to melt ice (, “enthalpy of fusion”) equals 6.01 kJ/mol and the enthalpy to boil water (, “enthalpy of vaporization”) equals 40.7 kJ/mol.

Pre-class question 1: Consider intermolecular forces and state why . Then, using section 8.7, calculate the value for , the enthalpy required to convert ice directly into vapor.

Fact: The H for the reaction, PCl5(g) PCl3(g) + Cl2(g), is +87.9 kJ/mol. This indicates that heat is going into the reactants to make the reaction occur.

Pre-Class Question 2: Count the number of bonds broken and bonds made in this reaction and state whether there are more bonds made or broken.

Fact: The H for the reaction, N2(g) + 3 H2(g) 2 NH3(g), is –92.2 kJ/mol. This indicates that heat is going out of the reactants to make the reaction occur.

Pre-Class Question 3: Count the number of bonds broken and bonds made in this reaction and state whether there are more bonds made or broken.

Pre-Class Question 4: Answer the following summary question: when there is (choose one: more/less) bond breaking than making. And when there is (choose one: more/less) bond breaking than making.

Check yourself: Do problem 8.10 (after reading the associated section) and explain why one reaction is endothermic and the other is exothermic. (The answers to all problems within the chapter are found at the end of the book.)

Check again: Do problem 8.11 and check the answer. Write down your strategy in your own words. Looking ahead to section 8.5, which would this reaction heat up more, 100 mL of water or 200 mL of water?

Application question 1: Work through Example 8.5. Explain why the reaction of Al(s) with Fe2O3(s) is exothermic (hugely, see ) while the “opposite” reaction, Fe(s) with Al2O3(s), is endothermic (boring). Estimate the mass of Fe2O3(s) needed to fill a flower pot and state the total energy given off by the reaction.

Application question 2: In the thermite video above, they blow up 8 gallons of gas. Assume that 8 gallons of gas weighs 23 kg and calculate the amount of energy released by the following reaction (gasoline is a complex mix of hydrocarbons, but we’ll assume it’s all octane for simplicity): C8H18() + O2 8 CO2(g) + 9 H2O(l), , where means “heat of combustion.”

In-class discussion:

Consider the following reactions. Be able to explain the chemical reason for the sign (+ or -):

Hot packs, exothermic: CaCl2(s) Ca2+(aq) + 2 Cl-(aq) H° = -81.35 kJ/mol.

Just plain hot -watch out! NaOH(s) Na+(aq) + OH-(aq) H° = -44.50 kJ/mol.

Cold packs, endothermic: NaCH3CO23H2O(s) Na+(aq) + CH3CO2-(aq) + 3H2O(l)

H° = +19.66 kJ/mol.

Another cold pack: NH4NO3(s) NH4+(aq) + NO3-(aq) H° = +25.7 kJ/mol.

Combustion reactions; what’s the advantage of methane over coal as a fuel? Section 8.11.

Methane: CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) H° = -802.34 kJ/mol.

Coal: C(s) + O2(g) CO2(g) H° = -393.51 kJ/mol.

Hydrogen: H2(g) + ½ O2(g) H2O(g) H° = -241.82 kJ/mol.

Why are there so many efforts to harness nuclear fusion?

H° = -1×109 kJ/mol!

Homework for section 8.7: See Mastering Chemistry for video tutorials; 8.64, 8.66, 8.68, due 9/1/23.

Topic 2, 9/1/23 (Section 8.8): how do we calculate enthalpy of reactions from other known reactions?

Hess’s Law: Enthalpy is a state function, which means that it measures the difference in energy between products and reactants, but it doesn’t consider how the reaction occurred. (Read sections 8.2 and 8.8 to fully understand that sentence.) The result is that we can combine any set of reactions to determine the enthalpy of a new reaction, just like adding algebraic equations. This is called Hess’s Law.

Pre-Class Question 5: Work through Examples 8.6 and 8.7 and be sure you can calculate (on your own) the for the combustion of methane (CH4) and the for the generation of H2 from C(s) and H2O(g). Notice that both reactions generate CO2 – why is that a concern?

Pre-Class Question 6: Do problems 8.12 – 8.14 and state whether you could draw equivalent arrow diagrams for Examples 8.6 and 8.7.

In-class discussion:

Let’s start off with an easy example involving an important reaction, the dissolution of limestone by acid rain: 2 H+(aq) + CaCO3(s) Ca2+(aq) + H2O(l) + CO2(g). (Google this issue).

Add the reactions below together like algebra equations, crossing out compounds that appear on both sides. Then add together the rH’s to determine the overall rH. The number is not huge, but it is negative. Why is that?

Dissociation of limestone: CaCO3(s) Ca2+(aq) + CO32-(aq) = -12.3 kJ/mol.

Carbonic acid formation: 2 H+(aq) + CO32-(aq) H2CO3(aq) = -22.6 kJ/mol.

Dissociation of carbonic acid: H2CO3(aq) H2O(l) + CO2(g) = +20.4 kJ/mol.

Show off your new-found talent by doing Problem 8.122 from the textbook:

Hess’s law can be used to calculate reaction enthalpies for hypothetical processes that can’t be carried out in the laboratory. Set up a Hess’s law cycle that will let you calculate ΔrH° for the conversion of methane to ethylene:

You can use the following information:

1. ,

2. ,

3. ,

4. ,

Homework: See Mastering Chemistry for interactive tutorial; 8.30, 8.72, 8.78, 8.112. Due 9/6/23.

Topic 3, 9/6/23 (section 8.9, 8.11 too): Now let’s calculate reaction enthalpies from standard enthalpies of formation ().

Fact: In general, one can calculate the from the individual of the reactants and products: H(reaction) = sum of Hf(products) – sum of Hf(reactants). The enthalpy of formation () is defined as the enthalpy to form each reactant and product from the elements in their standard state (i.e., liquid, solid or gas.)

Handy example from reaction 3 above: For the reaction , the formation reactions for each reactant and product are:

2 C(s) + 2 H2(g) C2H4(g)

H2(g) H2(g)

2 C(s) + 3 H2(g) C2H6(g)

For which reaction above does ? Look up the values for C2H4(g) and C2H6(g) in Appendix B and calculate the overall for C2H4(g) + H2(g) C2H6(g) and compare it to the value reported above.

Pre-Class Question 7: Work through Examples 8.8 and 8.9 and be sure you can calculate (on your own) the values for the reactions listed in this section. Is the formation of CaO(s) exo- or endothermic? Why does cement generate heat when concrete sets? (Ask prof Google.) For more practice, check the values for the other reactions in this worksheet.

Pre-Class Question 8: Do problems 8.15 and 8.16 and, for each reaction, state whether it’s endo- or exothermic. Look at the nature of each reaction and give a physical explanation for the sign of in each case.

In-class discussion: Let’s verify the reaction enthalpies in Topic 1 and compare different fuels.

Combustion reactions; what’s the advantage of methane over coal as a fuel? Calculate for the following reactions from the of each reactant and product. Most of these compounds should be in Appendix B; ask prof Google for the others. Read section 8.11.

Methane: CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g) = ______ kJ/mol

Coal: C(s) + O2(g) CO2(g) = ______ kJ/mol.

Hydrogen: H2(g) + ½ O2(g) H2O(g) = ______ kJ/mol.

Glucose: C6H12O6(s) + 6O2 6 H2O(l) + 6 CO2(g) = ______ kJ/mol.

Butane: C4H10() + 19/2 O2 7 H2O(l) + 6 CO2(g) = ______ kJ/mol

The group members should divide up the above reactions and determine for each one.

Rank these fuels from most efficient to least. Be sure you clearly define what you mean by “efficient.”

Homework: 8.76, 80, 82, 84, 86. Due 9/8/23.

Topic 4, 9/8/23 (section 8.5): How do we measure reaction enthalpies experimentally?

Fact: The molar heat capacity (“C”) is the amount of energy required to increase the temperature of 1 mol of a substance by 1°C; it has units of . The specific heat capacity is the amount of energy required to increase the temperature of 1 g of a substance by 1°C; it has units of . Compare the molar and specific heat capacities for liquid vs. solid H2O in Table 8.1 and state which phase (water or ice) has a higher heat capacity and why. This will involve you waving your hands in the air.

Cool and useful Fact: When you carry out a reaction in water, you can use the temperature increase (or decrease) to determine the enthalpy of reaction. Read section 8.5 to see how this is done. In short, you equate the heat absorbed by the water to the heat given off by the reaction.

Pre-class question 9: Follow example 8.3 to determine the for the reaction of Ag+ with Cl-. Once you verify their result, calculate the temperature rise if you mix 20.0 mL of 1.00 M Ag+ with 20.0 mL of 1.00 M Cl-.

Pre-class question 10: do Problem 8.7 and write down the strategy you followed in your own words to solve this problem.

In-class discussion: Let’s apply these ideas to hot and cold packs.

Hot packs. All reactions are spontaneous; note that some are exothermic and some are endothermic. How can a spontaneous reaction increase in energy?

Hot packs, exothermic: CaCl2(s) Ca2+(aq) + 2 Cl-(aq) = -81.35 kJ/mol.

MgSO4(s) Mg2+(aq) + SO42-(aq) = -91.22 kJ/mol

Just plain hot! NaOH(s) Na+(aq) + OH-(aq) = -44.50 kJ/mol.

Pick one of the above reactions and calculate how much 1.0 g of the solid would increase the temperature of 100 mL of water. Let your group mates do the other ones.

Give a chemical explanation why these reactions are exothermic. Muster up your intuition and speculate why they are spontaneous.

Cold packs. All reactions are spontaneous; note that some are exothermic and some are endothermic. How can a spontaneous reaction increase energy (go uphill)?

Cold packs, endothermic: NH4NO3(s) NH4+(aq) + NO3-(aq) = +25.7 kJ/mol

NaCH3CO23H2O(s) Na+(aq) + CH3CO2-(aq) + 3H2O(l) = +19.66 kJ/mol

The reaction above is from Chemistry Stack Exchange, not in any textbook I own! In words, the solid reactant is “sodium acetate trihydrate,” which means the solid crystal actually incorporates individual molecules of H2O – not H2O(l).

Pick one of the above reactions and calculate how much 1.0 g of the solid would decrease the temperature of 100 mL of water. Let your group mates do the other one.

Give a chemical explanation why these reactions are endothermic. Muster up your intuition and speculate why they are spontaneous.

Homework: See Mastering Chemistry for Calorimetry Tutorial; 8.46, 48, 50, 52, 53, 124. Due 9/11/23.

Multi-concept question: Each group member should pick one of the fuels from three combustion reactions given at the end of page 2 above. Calculate how much of the fuel is needed to increase the temperature of 100.0 mL of water from 25°C to 100°C and then boil it (see section 8.7).

Challenge Problem: In summer, 2023, the ocean temperature was 4°C higher than normal for this time of year (actually more in some parts). Look up the volume of the Atlantic Ocean (where most hurricanes occur) and calculate the amount of heat energy required to raise its temperature by 4°C. Compare that to the amount of energy used worldwide in a year, 606 EJ, which equals 6.06×1020 J.

Group Report: Reaction Enthalpy, Ch. 8

Turn in one set of responses for the entire group. This is worth 6 points and is due by Wednesday, 9/13/23, but early submissions will get EC!

Pick A or B below and answer it fully.

Be sure your group members can do the Multi-Concept Question on P. 6 of the Ch. 8 notes regarding combustion reactions.  Which fuel heats water the most? How much of that fuel is needed to increase the temperature of 100.0 mL of water from 25°C to 100°C and then boil it?

Answer the Challenge problem at the end of the Ch. 8 notes: In summer, 2023, the ocean temperature was 4°C higher than normal for this time of year (actually more in some parts). Look up the volume of the Atlantic Ocean (where most hurricanes occur) and calculate the amount of heat energy required to raise its temperature by 4°C. Compare that to the amount of energy used worldwide in a year, 606 EJ, which equals 6.06×1020 J.

There are a number of reactions in items 1, 2 and 3 in the notes where is given to you. Calculate the for 4 of those reactions using ’s (enthalpies of formation) of reactants and products.

List the names of the group members and the role they played to ensure the group was successful in accomplishing their goals. List one strength of the group, one way the group can improve and one thing the group learned that they wouldn’t have if everyone worked individually.

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