Thermochemistry

1. Thermochemistry The Study of Energy in Chemical Reactions

2. EVERY Chemical Change Has an Associated Change of Energy • Name some specific reactions or types of reactions. Go ahead, just shout some out. Is there an energy change? • What are some types of energy you know?

3. Thermochemistry • The study of energy and its transformations is known as THERMODYNAMICS. • Thermodynamics studies the relationships between heat, work and the energy content of a system. • If a system has ENERGY, it has the capacity to “do work or transfer heat.”

4. Potential Energy Kinetic Energy Types of Energy—The “Big Two” • Kinetic Energy is the energy of MOTION. • Potential Energyis the energy of POSITION. • Potential Energy is NOThow much energy an object “potentially could have.”

5. A Note About Potential Energy Potential Energy is NOT how much energy something “potentially has.” Potential energy is an actual quantity that objects possess. Potential energy is stored energy. It can be “taken out of storage” if it is converted into kinetic energy. Do not confuse the traditional meaning of the word “potential” with the way it is used in “potential energy”! This is a pitfall for many students! To put it another way: I can put a couch in storage. If I do this, it will not be in use. However, that does NOT mean it will stop being a couch. Potential energy may not be in useas kinetic energy, but it is still energy. The energy exists and is contained in the object regardless of whether the object is in use or not.

6. Practice Problems Identify the Energy Types A. A gas is contained in a balloon. • Where is the kinetic energy? • Where is the potential energy? B. A solid salt is stirred into water until it dissolves. • Where is the kinetic energy? • Where is the potential energy?

7. Chemical Energy • Substances contain chemical energy that comes from the potential energy stored in the arrangements of their atoms. • Chemical substances have kinetic energyassociated with their temperature, which is caused by the movements of atoms and subatomic particles.

8. Units of Energy—SI • The SI unit for energy is the Joule. • in honor of James Joule, a British scientist who investigated work and heat. • Joules are a small unit, so kilojoules (kJ)are often used when discussing the energies associated with chemical reactions. • 1 kJ = 1000 J

9. Units of Energy—American • Calories (abbreviated as “Cal”) • A more familiar unit to us. • Defined as the amount of energy needed to raise the temperature of 1 gram of water from 14.5 °C to 15.5 °C. • 1 calorie = 4.184 J • 1 Calorie = 1 kcal = 1000 calories. NOTICE: THERE IS A DIFFERENCE BETWEEN “calories” and “Calories”—a capital “C” makes it a bigger unit!!

10. Practice Problems • Convert 8.42 J to calories. • Convert 1200 calories to Calories. • Convert 1.45 kcal to J.

11. Systems vs. Surroundings • The universe is divided into two halves. • System: the reaction or process actually being studied. • Surroundings: literally everything else—the container, the room, the people standing next to it, puppies, outer space, etc.

12. SURROUNDINGS SYSTEM • We only care about the system for the purposes of all chemical calculations. • Anything that is not part of the system—anything we don’t care about—is the surroundings.

13. Practice Problems • Gas is contained in a balloon. • What is the system? • What are the surroundings? • A solid salt is stirred into water until it dissolves. • What is the system? • What are the surroundings?

14. Exergonic vs. Endergonic • Exergonic reactions release energy to the surroundings. • Endergonic reactions absorb energy from the surroundings. • Which is exergonic and which is endergonic? exergonic endergonic

15. Practice Problems Tell whether the situation is endergonic or exergonic. • A plant absorbs sunlight to use in photosynthesis. • An engine creates heat and light. • A windmill rotates in the wind. • The rotation of a windmill turns a turbine in a power plant.

16. Exit Slip Questions • On your own notebook paper, please answer the following questions and turn them in to the folder on your way out of class. • What is thermochemistry? • How does potential energy usually factor into chemical situations? • What is the SI unit of energy? What is the American unit of energy?

17. The First Law of Thermodynamics The energy of the universe is constant. It may be transferred between the system and the surroundings, but overall, the amount of energy in the universe will not and can not change. This law is also called the Law of Conservation of Energy.

18. The First Law of Thermodynamics The short version: Energy is neither created, nor destroyed. Michael Phelp’s Breakfast

19. Specific Heat Capacity • Specific heat capacity (specific heat) is the heat required to raise the temperature of one gram of a substance to be 1 °C hotter. • We use the variable “c” (lower case!) • Constant value for a given substance—you can look it up in a table. • Unitsof J/g•°C

20. Specific Heat Capacity • We can determine how much heat is gained or lost to a system if we know its specific heat. Two other necessary pieces of information: • Mass • Temperature

21. q= mcΔT • q = heat (J) • m = mass (g) • c = specific heat capacity (J/g•°C) (look it up in a table, or it will be given to you) • ΔT = change in temperature (Final temperature minus initial temperature)

22. Solving q = mcΔT • Check that everythingis in the correctunits. Convert if necessary. (q in J, m in g, c in J/g•°C, T or ΔT in °C). • Look up the specific heat in a chart if it is not given to you. • If you are given Tfinal and Tinitialinstead of ΔT, subtract to find ΔTΔT = Tfinal – Tinitial • Plug everything in! Do algebra as necessary to solve the problem. Make sure you end up in the correct units, or you may have done something wrong! NOTICE that heat, q, is not solely dependent on temperature.Heat and temperature are not the same thing.

23. Practice Problems A. If 25.0 g of Al cool from 310 oC to 37 oC, how many Joules of heat energy are lost by the Al? q = m = c = ΔT = 6142.5 J 25.0 g 0.900 J/goC (see table on p. 4 of notes!) 37 – 310 = -273 oC q= mcΔT

24. Practice Problems B. What is the specific heat of a substance if 5684 J are needed to heat 123.4 grams by 12.4 °C? q = m = c = ΔT = 5684 J 123.4 g 3.715 J/goC 12.4 oC q= mcΔT

25. Practice Problems C. What mass of glass can be heated by 45.5 °C with 850.0 J of energy? q = m = c = ΔT = q= mcΔT 850.0 J 22.3 g 0.837 J/goC (See table on p. 4 of notes!) 45.5 oC

26. Practice Problems D. What will be the temperature change if 267.4 cal are added to 62.54 grams of silver? q = m = c = ΔT = 1 calorie = 4.18 Joules 267.4 cal = 1118.8 J 62.54 g 0.236 J/goC (See table!) 75.8 oC

27. Calorimetry • Heat flow—The temperature change experienced by an object when it absorbs a certain amount of heat. • Calorimetry is a procedure to measure heat flow. • A big specific heat means that more heat is required to raise an object’s temperature.

28. How Does Calorimetry Work? • If we’re trying to measure heat, can we do this directly? • q = mcΔT says we need a mass, a specific heat, and a temperature change. • We know that any heat created by a reaction can be transferredfrom the reaction’s system to its surroundings. • We know that any heat gained by the surroundings MUST come from the system. • How can we use this knowledge to set up an experimental situation where we can measure this heat transfer?

29. How Does Calorimetry Work? To measure something’s heat change, we need to measure the mass and the change in temperature. We also need to know the specific heat capacity of our object. • If I heated up water, how could I find the heat change? Remember q = mcΔT! • Measure the water’s mass with a scale and its change in temperature with a thermometer. Look up the specific heat of water in a table. Plug in to q = mcΔT.

30. How Does Calorimetry Work? To measure something’s heat change, we need to measure the mass and the change in temperature. We also need to know the specific heat capacity of our object. • If I heated up a penny, how could I find the heat change? • Problem! A thermometer will not really work well with a penny. How can you find the penny’s temperature change?

31. Your turn: Invent Calorimetry In your groups, discuss and produce a procedure for a laboratory experiment that would measure the heat change from a system to its immediate surroundings. Think about these two questions: • If I heated up some water, how could I measure its heat change, q? • If I heated up a copper penny, how could I measure its heat change, q? Remember: q = mcΔT! How will you measure mass? How will you measure change in temperature? Where will you get “c” from?

32. Class discussion • We can’t use a thermometer to measure a penny’s temperature change…but we can measure water’s! • If your group realized that you could put the hot penny into some water and measure how much the water’s temperature changed, congratulations! You invented calorimetry.

33. A procedure for calorimetry • Measure the mass of the water you will be using in your calorimeter. • Put a thermometer in the water. Measure its initial temperature. • Add whatever object or reaction you want to measure the temperature change of.

34. A procedure for calorimetry (continued) 4. Quickly close the calorimeter. 5. Watch the temperature and record how high (or how low!) it gets. 6. Calculate the heat change of the water using q = mcΔT for water. All the heat gained by the water had to be lost by the object or reaction. This means that the “q” for water is the same amount of heat change for your object!

35. Endothermic = absorbing energy Surroundings everything else Exothermic = releasing energy System with can as boundary A basic calorimeter Law of conservation of energy = release and absorption of energy must be equal

36. Bomb Calorimetry • Calorimetry that takes place in a closed volume. http://www.youtube.com/watch?v=ohyA9amFfsc

37. Enthalpy—ΔH • From a Greek word meaning “to warm” • Accounts for heat flow in processes occurring at constant pressure. • We can only measure how much enthalpy changes from one moment to the next.

38. Enthalpy—ΔH • At constant pressure (and only at constant pressure), ΔH = q (heat). • The change in enthalpy equals the heat gained or lost at constant pressure.

39. Endothermic vs. Exothermic • A positive value of ΔH means that enthalpy (heat) is entering the system from the surroundings. Reactions like this are endothermic. • A negative value of ΔH means that enthalpy (heat) is leaving the system and moving into the surroundings. Reactions like this are exothermic.

40. Enthalpy Values • Every chemical reaction involves an exchange of energy • Heat flow from the system to surroundings • Exothermic • Heat flow from the surroundings to system • Endothermic

41. Enthalpy Values • We always view enthalpy change from the point of view of the system. This means that the sign of enthalpy is positive when the system gains heat energy, and the sign of enthalpy is negative when the system loses heat energy.

42. I need four volunteer actors, please.

44. Hess’s Law The formula: ΔH°reaction = ΣΔHf°(products) – ΣΔHf°(reactants) Enthalpy change for the overall reaction Minus The sum of the enthalpy change for all the reactants together (find them in a table) The sum of the enthalpy change for all the products together (find them in a table)

45. Hess’s Law In words: The enthalpy change of a reaction equals the sum of the enthalpy change of the products minus the sum of the enthalpy change of the reactants.

46. Practice Problems • What is the enthalpy change for the following reaction? Use the table of enthalpy values in your notes. Is this reaction endothermic or exothermic? BaCO3 + FeSO4-----> BaSO4 + FeCO3 ΔHf° = -1213 kJ/mol ΔHf° = -1473.2 kJ/mol ΔHf° = -929 kJ/mol ΔHf° = -750.6 kJ/mol ΔH°reaction= (-1473.2 + -750.6) – (-1213 + -929) ΔH°reaction= (-2223.8) – (-2142) ΔH°reaction = -81.8 kJ/mol ΔH°reaction is NEGATIVE. This means the system LOSES heat. If heat EXITS the system, the reaction is EXOthermic.

47. Practice Problems B. What is the enthalpy change for the following reaction? Use the table of enthalpy values in your notes. Is this reaction endothermic or exothermic? CdSO4 + MgO-----> CdO + MgSO4 ΔHf° = -935 kJ/mol ΔHf° = -258 kJ/mol ΔHf° = -601.24 kJ/mol ΔHf° = -1278.2 kJ/mol ΔH°reaction= (-258 + -1278.2 ) – (-935 + -601.24) ΔH°reaction= (-1536.2) – (-1536.24) ΔH°reaction = 0.04 kJ/mol ΔH°reaction is POSITIVE. This means the system GAINS heat. If heat ENTERS the system, the reaction is ENDOthermic.

48. Practice Problems C. What is the enthalpy change for the following reaction? Use the table of enthalpy values in your notes. Is this reaction endothermic or exothermic? Ca + Cl2-----> CaCl2 ΔHf° = 0 kJ/mol ΔHf° = -877.3 kJ/mol ΔHf° = 0 kJ/mol ΔH°reaction= (-877.3) – (0) ΔH°reaction = -877.3 kJ/mol A negative enthalpy value means the reaction is exothermic!

49. Enthalpy Values for Elemental States Did you notice that some of the values in the chart were zero? What do all the species with ΔH = 0 have in common? • They are all in their “elemental state.” This means that the element is the way it occurs in nature. It takes no energy for these elements to be in their natural state, so there is no heat flow to create them.

50. Elemental States For most elements, the Elemental State is just one atom of the element. For diatomic elements, the elemental state is two atoms together in one molecule. Which elements are we talking about here? Br2 I2 N2 Cl2 H2 O2 F2 !!!!