1. Thermodynamics

2. Energy

3. Energy in general is the ability to cause a change. In chemistry, energy can do work or produce heat. Energy is typically divided into two types: potential energy and kinetic energy.

4. Potential energy is stored energy in the form of position (gravity-related, like a skateboarder at the top of a half-pipe) or chemical composition (like the chemicals in food or a battery).

5. Kinetic energy is energy of motion, and is the usual use for potential energy (food energy allows you to move). A chemical reaction could have both kinetic energy, as the atoms or molecules move around, and potential energy, as there could be chemicals still waiting to react.

6. Energy can be converted between kinetic energy and potential energy, even several times, but no matter how many times energy is converted the Law of Conservation of Energy states that no energy can be lost or created.

7. The energy stored in a substance because of the atoms that make it up is chemical potential energy. For example, when octane (C 8 H 18 ) is placed into an automobile’s gas tank, it can be burned later to move the automobile. Unfortunately a lot of the chemical potential energy in gasoline is “lost” in the engine as heat, so it is not available for motion. Heat

8. Heat Energy

9. Heat (symbolized q) is the energy that naturally moves from something hot to something cold. The amount of heat is often measured as temperature, but temperature does not directly measure heat. For example, water at 20 °C does not have twice as much heat as water at 10 °C, but it does have more heat!

10. To measure heat directly, the calorie (abbreviated cal) was defined as the amount of heat needed to increase the temperature of 1 gram (1 milliliter) of water by 1 °C. Note that the Calories reported on food wrappers are the measurement for heat to increase the temperature of 1000 grams (one liter) of water by 1 °C, in other words a kilocalorie.

11. Drinking a can of Pepsi actually means drinking 150,000 calories! Keep in mind that a normal person’s diet is probably around 2,000,000 calories per day. The measurement of how much heat energy can come from food is called calorimetry.

12. Calories are still used in the nutrition field, but in chemistry we will use the Joule (abbreviated J). There are 4.184 Joules in 1 calorie, so Joules can measure energy a little more accurately. Both the calorie and Joule are defined using pure water, but only pure water gains and loses heat like pure water. Everything else in the world gains and loses heat at a different rate – a factor called the specific heat (abbreviated c).

13. When a cake has been baked and is removed from the oven it would cause severe burns to touch the metal pan, but often the top of the cake is touched to test if the cake is ready. The metal of the pan has a low specific heat so it can lose heat quickly (which would be absorbed by your finger) and the cake contains a lot of trapped air and other material with a higher specific heat which lose heat slowly, allowing more time before your finger has absorbed enough heat to cause a burn.

14. Additionally, the density of the cake is usually much less than the density of the metal pan, so for the same surface area (the surface area touched by your finger) there are a lot more metal atoms that could transfer heat energy than cake molecules. Thus not only does the amount of heat that could be transferred relate to the specific heat, the mass also has an affect. (Also, metals tend to be good conductors of heat while air is a poor conductor of heat.)

15. In total, to calculate the amount of heat that can be gained or lost by an object the formula is q = m × c × ∆T q = heat, measured in J m = mass, measured in g c = specific heat, measured in ∆T = Temperature final – Temperature initial, measure in °C J g °C

16. If the object is heating up, then the final temperature is greater than the initial temperature, so ∆T is positive and the number for heat is positive. If the object is cooling down, then the final temperature is less than the initial temperature, so ∆T is negative and the number for heat is negative. - q+ q

17. What is the heat absorbed by 100 g of water to raise it from 25 °C to 75 °C? The specific heat of water is 4.184 J/g°C. q = m × c × ∆T q = ________ x ________ x (________ - _______) q = _______________

18. What is the heat given off by 100 g of iron to cool from 50 °C to 20 °C? The specific heat of iron is 0.449 J/g°C. q = m × c × ∆T q = ________ x ________ x (________ - _______) q = _______________

19. Heat Movement

20. The Law of Conservation of Energy reminds us that heat energy cannot just appear out of nowhere and disappear to nowhere, so if an object warms up, then that heat had to be absorbed from somewhere and if an object cools down, then that heat had to be given off to something.

21. In our world ecosystem heat is gained by the sun’s radiation and given off to space at night. This is why a cloudy night is typically warmer than a clear night – the heat gets trapped in by the clouds.

23. Heat can move from one object to another by three ways: convection, conduction, and radiation.

25. It was already mentioned how the energy from the sun reaches earth by radiation, but radiant heat energy is the movement of heat by electromagnetic waves.

26. The infrared portion of the electromagnetic spectrum is the “heat” form of radiation. Other examples of radiation are the heat from a fire, non- fluorescent lightbulbs, and the coil in an electric oven that heats up to cook the food.

27. Convection is the movement of heat by circulating materials, so anything that is a fluid (gases and liquids) can move heat by convection, often called convection currents. Examples of convection include most air- conditioning systems, the ocean currents, and even the warm and cold “fronts” the weather-person talks about on the news.

29. Conduction is the movement of heat between objects that are touching. This is the most common form of heat movement when cooking. The food touches a pan that is touching a burner, and as the pan absorbs heat from burner the food absorbs heat from the pan. When you touch a hot pan heat moves from the pan into your hand and gives you a burn. Often the lasting effects of a burn can be minimized if you can quickly get the heat to move out of your hand by running cold water over your hand for a long time.

30. Regardless of the method of heat transfer, heat energy will always naturally move from an object with higher heat to an object with lower heat. Also, the Law of Conservation of Energy tells us that the amount of heat lost by the hot object must be equal to the amount of heat gained by the cold object. As heat lost is a negative number, we will need to add a negative sign to the lost heat number so that it matches the gained heat number.

31. Hot becomes cold, and cold becomes hot. - q lost = q gained so - (m lost × c lost × ∆T lost ) = m gained × c gained × ∆T gained

32. Phases and Changes

33. Sometimes when an object gains or loses heat, it is not just the temperature that changes but it could be that the phase of the object changes. The phases we are concerned with are solid, liquid, and gas (vapor) - which are sometimes called the states of matter. The phase that an object currently exists in depends on how much heat energy is in the object, often measured by the temperature. Thus the solid form of any object is going to have lower temperature than the liquid form, which will have lower temperature than the gas (vapor) form.

34. Phase changes occur when sufficient heat energy is added or removed to change the object’s phase. Melting is the solid phase changing to the liquid phase, and freezing is the liquid phase turning to a solid phase. Melting and freezing both occur at the melting point – the temperature at which a solid becomes a liquid. If heat is being absorbed (+q), then melting is occurring. If heat is given off (-q), then freezing is happening.

35. Boiling (vaporizing) is the liquid phase changing to the gas phase, and condensation is the gas phase changing to the liquid phase. Boiling and condensing occur at the boiling point – the temperature at which a liquid becomes a gas. If heat is being absorbed (+q), then boiling is occurring. If heat is given off (-q), then condensing is occurring. Sometimes liquids evaporate instead of boiling.

36. Evaporation is boiling, but occurs below the boiling point. Evaporation can only happen at the surface of the liquid (boiling happens everywhere inside) and is a cooling process.

37. There are two other phase change possibilities, but they are rare in everyday life. Sublimation is the solid phase changing directly to the gas phase (no liquid in-between), and deposition is the gas phase changing directly to the solid phase. Once again, a +q would indicate sublimation and a –q would indicate deposition.

38. Pressure also has a small effect on the phases. For a liquid to boil, the pressure inside the liquid needs enough energy to equal the pressure outside the liquid. In Denver where there is less atmospheric pressure than here in Haltom City, the water does not need as much heat energy to boil, so water will boil around 95 °C. Most of the time, the solid phase is more dense than a liquid, so high pressure can force a liquid to become a solid, even if the temperature is above the normal melting point.

39. To place all this information about the pressure, temperature, and phases into a simple to read form, chemists construct a phase diagram.

40. Solid Liquid Gas

41. Solid Liquid GasSublimation Melting Boiling

42. Solid Liquid GasSublimation Melting Boiling Melting PointBoiling Point

43. The phase diagram is helpful to compare the phase that will exist at any particular pressure and temperature, but if the amount of energy that is needed to change a phases’ temperature or change the phase is needed, it is easier to put the information into a heat curve.

44. From this heat curve (for water) we can see how the temperatures change and the phases change. Notice that the temperature does not change during a phase change (the two flat lines). This means our equation for heat q = m × c × ∆T will not work during a phase change, but it makes sense that heat is needed to melt something

45. so new equations are needed for use during a phase change: during melting/freezing, q = m × H f q = heat (add “-“ sign for freezing), measured in J m = mass, measured in g H f = heat of fusion, measured in J g

46. so new equations are needed for use during a phase change: during boiling/condensing, q = m × H v q = heat (add “-“ sign for condensing), measured in J m = mass, measured in g H v = heat of vaporization, measured in J g

47. Example Calculations Calculate the heat needed for the heat curve above for 100 g of water. H f water = 334 J/g H v water = 2260 J/g c ice = 2.06 J/g °C c water = 4.18 J/g °C c steam = 2.08 J/g °C

48. Step 1: Ice from -40 °C to 0 °C. q = m × c × ∆T q = ________ x ________ x (________ - _______) q = _______________

49. Step 2: Turning ice into liquid water q = m × Hf q = ___________ x ___________ q = _______________

50. Step 3: Water from 0 °C to 100 °C. q = m × c × ∆T q = ________ x ________ x (________ - _______) q = _______________

51. Step 4: Turning liquid water into water vapor (steam). q = m × Hv q = ___________ x ___________ q = _______________

52. Step 5: Water vapor from 100 °C to 140 °C q = m × c × ∆T q = ________ x ________ x (________ - _______) q = _______________

53. Total heat involved = ________ + _______ + _______ + _______ + _______ Total heat = __________

54. Heat and Chemical Reactions

55. Although we did not write heat as a reactant or product back in chapter 10, every chemical reaction involves a change in energy. If a chemical reaction feels cold when it is happening, then the reaction is absorbing heat. If a chemical reaction feels warm when it is happening, then it is giving off heat.

56. When energy is absorbed by a chemical reaction, it is written as a reactant, and is called endothermic. For example: solid + energy (heat) → liquid For this reaction to happen, heat needs to be removed from the environment surrounding the reaction for it to occur. This is why your skin gets cold when you hold ice; the ice is pulling heat out of your skin to melt.

57. When energy is given off by a chemical reaction, it is written as a product, and is called exothermic. For example: Iron + oxygen gas → iron(III) oxide + energy (heat) When this reaction happens heat is given off by the chemical reaction, which is why it is used in instant hot pads.

58. In both endothermic and exothermic chemical reactions, the energy does not need to be heat only. Some chemicals are stored in dark brown bottles to block out the light. Light can cause some chemicals to deteriorate faster than a chemist would like, ruining the purity of the chemical. To separate water into hydrogen and oxygen (hydrolysis) electricity is often used. Is hydrolysis of water endothermic or exothermic?

59. Some chemical reactions give off light, like the reaction in a glow stick. When hydrogen peroxide reacts with a phenyl oxalate ester energy is given off, but in the form of light and only a little heat.

60. Sometimes the energy in a chemical reaction can be used to speed up or even reverse the reaction. Most chemical reactions are reversible and usually both the forward and reverse reactions are occurring at the same time, but at the same rate so you will not notice a difference, called an equilibrium reaction. Changing the energy can remove the equilibrium and then a change is noticed, like ice melting in a cup sitting on the counter and freezing when the cup is moved into a freezer. ice + heat ↔ liquid water