Chapter 11 - Thermochemistry Heat and Chemical Change

1. Chapter 11 - ThermochemistryHeat and Chemical Change

2. Section 11.1The Flow of Energy - Heat • OBJECTIVES: • Explain the relationship between energy and heat.

3. Section 11.1The Flow of Energy - Heat • OBJECTIVES: • Distinguish between heat capacity and specific heat.

4. Energy and Heat • Thermochemistry - concerned with heat changes that occur during chemical reactions • Energy - capacity for doing work or supplying heat • weightless, odorless, tasteless • if within the chemical substances- called chemical potential energy

5. Energy and Heat • Gasoline contains a significant amount of chemical potential energy • Heat - represented by “q”, is energy that transfers from one object to another, because of a temperature difference between them. • only changescan be detected! • flows from warmer  cooler object

6. Exothermic and Endothermic Processes • Essentially all chemical reactions, and changes in physical state, involve either: • release of heat, or • absorption of heat

7. Exothermic and Endothermic Processes • In studying heat changes, think of defining these two parts: • the system - the part of the universe on which you focus your attention • the surroundings - includes everything else in the universe

8. Exothermic and Endothermic Processes • Together, the system and it’s surroundings constitute the universe • Thermochemistry is concerned with the flow of heat from the system to it’s surroundings, and vice-versa. • Figure 11.3, page 294

9. Exothermic and Endothermic Processes • The Law of Conservation of Energy states that in any chemical or physical process, energy is neither created nor destroyed. • All the energy is accounted for as work, stored energy, or heat.

10. Exothermic and Endothermic Processes • Fig. 11.3a, p.294 - heat flowing into a system from it’s surroundings: • defined as positive • q has a positive value • called endothermic • system gains heat as the surroundings cool down

11. Exothermic and Endothermic Processes • Fig. 11.3b, p.294 - heat flowing out of a system into it’s surroundings: • defined as negative • q has a negative value • called exothermic • system loses heat as the surroundings heat up

12. Exothermic and Endothermic • Fig. 11.4, page 295 - on the left, the system (the people) gain heat from it’s surroundings (the fire) • this is endothermic • On the right, the system (the body) cools as perspiration evaporates, and heat flows to the surroundings • this is exothermic

13. Exothemic and Endothermic • Every reaction has an energy change associated with it • Exothermic reactions release energy, usually in the form of heat. • Endothermic reactions absorb energy • Energy is stored in bonds between atoms

14. Heat Capacity and Specific Heat • A calorie is defined as the quantity of heat needed to raise the temperature of 1 g of pure water 1 oC. • Used except when referring to food • a Calorie, written with a capital C, always refers to the energy in food • 1 Calorie = 1 kilocalorie = 1000 cal.

15. Heat Capacity and Specific Heat • The calorie is also related to the joule, the SI unit of heat and energy • named after James Prescott Joule • 4.184 J = 1 cal • Heat Capacity - the amount of heat needed to increase the temperature of an object exactly 1 oC

16. Heat Capacity and Specific Heat • Specific Heat Capacity - the amount of heat it takes to raise the temperature of 1 gram of the substance by 1 oC (abbreviated “C”) • often called simply “Specific Heat” • Note Table 11.2, page 296 • Water has a HUGE value, compared to other chemicals

17. Heat Capacity and Specific Heat • For water, C = 4.18 J/(g oC), and also C = 1.00 cal/(g oC) • Thus, for water: • it takes a long time to heat up, and • it takes a long time to cool off! • Water is used as a coolant! • Note Figure 11.7, page 297

18. Heat Capacity and Specific Heat • To calculate, use the formula: • q = mass (g) x T x C • heat abbreviated as “q” • T = change in temperature • C = Specific Heat • Units are either J/(g oC) or cal/(g oC) • Sample problem 11-1, page 299

19. Section 11.2Measuring and Expressing Heat Changes • OBJECTIVES: • Construct equations that show theheat changes for chemical and physical processes.

20. Section 11.2Measuring and Expressing Heat Changes • OBJECTIVES: • Calculate heat changes in chemical and physical processes.

21. Calorimetry • Calorimetry - the accurate and precise measurement of heat change for chemical and physical processes. • The device used to measure the absorption or release of heat in chemical or physical processes is called a Calorimeter

22. Calorimetry • Foam cups are excellent heat insulators, and are commonly used as simple calorimeters • Fig. 11.8, page 300 • For systems at constant pressure, the heat content is the same as a property called Enthalpy (H) of the system

23. Calorimetry • Changes in enthalpy = H • q = H These terms will be used interchangeably in this textbook • Thus, q = H = m x C x T • H is negative for an exothermic reaction • H is positive for an endothermic reaction (Note Table 11.3, p.301)

24. Calorimetry • Calorimetry experiments can be performed at a constant volume using a device called a “bomb calorimeter” - a closed system • Figure 11.9, page 301 • Sample 11-2, page 302

25. C + O2 Energy C O2 Reactants ® Products C + O2® CO2 + 395 kJ 395kJ

26. O C O O C O O C O In terms of bonds O C O Breaking this bond will require energy. Making these bonds gives you energy. In this case making the bonds gives you more energy than breaking them.

27. Exothermic • The products are lower in energy than the reactants • Releases energy

28. CaO + CO2 Energy CaCO3 Reactants ® Products CaCO3® CaO + CO2 CaCO3 + 176 kJ ® CaO + CO2 176 kJ

29. Endothermic • The products are higher in energy than the reactants • Absorbs energy • Note Figure 11.11, page 303

30. Chemistry Happens in MOLES • An equation that includes energy is called a thermochemical equation • CH4 + 2O2® CO2 + 2H2O + 802.2 kJ • 1 mole of CH4 releases 802.2 kJ of energy. • When you make 802.2 kJ you also make 2 moles of water

31. Thermochemical Equations • A heat of reaction is the heat change for the equation, exactly as written • The physical state of reactants and products must also be given. • Standard conditions for the reaction is 101.3 kPa (1 atm.) and 25 oC

32. CH4 + 2 O2® CO2 + 2 H2O + 802.2 kJ • If 10. 3 grams of CH4 are burned completely, how much heat will be produced? 1 mol CH4 802.2 kJ 10. 3 g CH4 16.05 g CH4 1 mol CH4 = 514 kJ

33. CH4 + 2 O2® CO2 + 2 H2O + 802.2 kJ • How many liters of O2 at STP would be required to produce 23 kJ of heat? • How many grams of water would be produced with 506 kJ of heat?

34. Summary, so far...

35. Enthalpy • The heat content a substance has at a given temperature and pressure • Can’t be measured directly because there is no set starting point • The reactants start with a heat content • The products end up with a heat content • So we can measure how much enthalpy changes

36. Enthalpy • Symbol is H • Change in enthalpy is DH (delta H) • If heat is released, the heat content of the products is lower DH is negative (exothermic) • If heat is absorbed, the heat content of the products is higher DH is positive (endothermic)

37. Energy Change is down DH is <0 Reactants ® Products

38. Energy Change is up DH is > 0 Reactants ® Products

39. Heat of Reaction • The heat that is released or absorbed in a chemical reaction • Equivalent to DH • C + O2(g) ® CO2(g) + 393.5 kJ • C + O2(g) ® CO2(g) DH = -393.5 kJ • In thermochemical equation, it is important to indicate the physical state • H2(g) + 1/2O2 (g)® H2O(g) DH = -241.8 kJ • H2(g) + 1/2O2 (g)® H2O(l) DH = -285.8 kJ

40. Heat of Combustion • The heat from the reaction that completely burns 1 mole of a substance • Note Table 11.4, page 305

41. Section 11.3Heat in Changes of State • OBJECTIVES: • Classify, by type, the heat changes that occur during melting, freezing, boiling, and condensing.

42. Section 11.3Heat in Changes of State • OBJECTIVES: • Calculate heat changes that occur during melting, freezing, boiling, and condensing.

43. Heats of Fusion and Solidification • Molar Heat of Fusion (Hfus) - the heat absorbed by one mole of a substance in melting from a solid to a liquid • Molar Heat of Solidification(Hsolid) - heat lost when one mole of liquid solidifies

44. Heats of Fusion and Solidification • Heat absorbed by a melting solid is equal to heat lost when a liquid solidifies • Thus, Hfus = -Hsolid • Note Table 11.5, page 308 • Sample Problem 11-4, page 309

45. Heats of Vaporization and Condensation • When liquids absorb heat at their boiling points, they become vapors. • Molar Heat of Vaporization(Hvap) - the amount of heat necessary to vaporize one mole of a given liquid. • Table 11.5, page 308

46. Heats of Vaporization and Condensation • Condensation is the opposite of vaporization. • Molar Heat of Condensation(Hcond) - amount of heat released when one mole of vapor condenses • Hvap = - Hcond

47. Heats of Vaporization and Condensation • Note Figure 11.5, page 310 • The large values for Hvapand Hcondare the reason hot vapors such as steam is very dangerous • You can receive a scalding burn from steam when the heat of condensation is released!

48. Heats of Vaporization and Condensation • H20(g) H20(l) Hcond = - 40.7kJ/mol • Sample Problem 11-5, page 311

49. Heat of Solution • Heat changes can also occur when a solute dissolves in a solvent. • Molar Heat of Solution (Hsoln) - heat change caused by dissolution of one mole of substance • Sodium hydroxide provides a good example of an exothermic molar heat of solution:

50. Heat of Solution NaOH(s)  Na1+(aq) + OH1-(aq) Hsoln = - 445.1 kJ/mol • The heat is released as the ions separate and interact with water, releasing 445.1 kJ of heat as Hsoln thus becoming so hot it steams! • Sample Problem 11-6, page 313 H2O(l)