Web viewUnit 3 Study Guide. Energy and Thermochemistry. 10.1 - Energy, Temperature, and Heat. 1....
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Unit 3 Study Guide Energy and Thermochemistry 10.1 - Energy, Temperature, and Heat 1. Explain the difference between kinetic and potential energy in a chemical reaction. 2 a. What do potential, kinetic and thermal energy have to do with the law of conservation of energy? b. What do they have to do with the internal energy of a system? 3. Explain the difference between temperature and heat. 4. Give an example of a state function and explain your answer. 5. Give an example of a system and explain your answer.
Transcript of Web viewUnit 3 Study Guide. Energy and Thermochemistry. 10.1 - Energy, Temperature, and Heat. 1....
Unit 3 Study GuideEnergy and Thermochemistry
10.1 - Energy, Temperature, and Heat
1. Explain the difference between kinetic and potential energy in a chemical reaction.
2 a. What do potential, kinetic and thermal energy have to do with the law of conservation of energy?
b. What do they have to do with the internal energy of a system?
3. Explain the difference between temperature and heat.
4. Give an example of a state function and explain your answer.
5. Give an example of a system and explain your answer.
10.2 - Endothermic vs. Exothermic
6. Draw the energy diagram for an endothermic and exothermic reaction:
7.a. Which has higher energy in an endothermic reaction, reactants or products?
b. Which has higher energy in an exothermic reaction, the reactants or products?
8. Label the process as either exothermic or endothermic and identify your perspective:
a. When solid KBr (potassium bromide) is dissolved in water, the solution gets colder ____________
b. Natural gas CH4 (methane) is burned in a furnace _______________
c. Concentrated H2SO4 (sulfuric acid) is added to water,and the solution gets very hot _____________________
d. Water boils in a teakettle _____________
10.3 - Thermodynamics
9. Explain what each positive or negative sign means:
State Function? Definition + positive - negative
Work (w)
Heat (q)
∆E = energy
∆H = enthalpy
∆T = temp
10. If a gas absorbs 45 kilojoules of heat, and does 29 kilojoules of work, calculate the change in energy.
11. Look at the specific heats listed in Table 10.1 in the textbook. Which substance retains the most heat? Which substance will have the largest change in temperature when heat is applied?
12. How much energy is required to heat 7.40 mL of water from 25°C to 46°C?
13. How much energy is required to heat a piece of iron (Fe) that weighs 1.3 grams from 25°C to 46°C? (Use table 10.1)
14. If it takes 5.8 joules of energy to heat a piece of metal that weighs1.6 grams from 23°C to 41°C, what is the specific heat of this metal? Is this metal pure gold? Why or why not? (Consult table 10.1)
15. How much energy is required to melt 23 g of ice?
16. Explain how a calorimeter could be used to solve question 15 experimentally. Would you expect the answer to be exactly the same as what you calculated in 15? Why or why not?
Thermochemistry (Enthalpy) and Hess’s Law
17. The equation for the fermentation of glucose to alcohol and carbon dioxide is:
C6H12O6 2 C2H5OH + 2 CO2 ∆H = - 67 kilojoules
a. Is this reaction endothermic or exothermic? Why?
b. Is energy released or absorbed?
c. How much heat is released when 25 moles of glucose is fermented?
18. S + 3/2 O2 SO3 ∆H = - 395.2 kilojoules
2SO2 + O2 2 SO3 ∆H = - 198.2 kilojoules
Calculate the Enthalpy (∆H) for the overall reaction:
S + O2 SO2
19. 2O3 3O2 ∆H = - 427 kilojoules
O2 2 O ∆H = + 459 kilojoules
NO + O3 NO2 + O2 ∆H = - 199 kilojoules
Calculate the Enthalpy (∆H) for the overall reaction:
NO + O NO2
