1 / 59

Chapter 11 - Thermochemistry Heat and Chemical Change

Chapter 11 - Thermochemistry Heat and Chemical Change. Section 11.1 The Flow of Energy - Heat. OBJECTIVES: Explain the relationship between energy and heat . Section 11.1 The Flow of Energy - Heat. OBJECTIVES: Distinguish between heat capacity and specific heat . Energy and Heat.

narcisse
Download Presentation

Chapter 11 - Thermochemistry Heat and Chemical Change

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  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)

More Related