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Chapter 11 - Thermochemistry Heat and Chemical Change

Chapter 11 - Thermochemistry Heat and Chemical Change. Milbank High School. 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 .

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Chapter 11 - Thermochemistry Heat and Chemical Change

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  1. Chapter 11 - ThermochemistryHeat and Chemical Change Milbank High School

  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

  5. Energy and Heat • Heat - represented by “q”, is energy that transfers from one object to another, because of a temperature difference between them.

  6. Exothermic and Endothermic Processes • In studying heat changes, think of defining these two parts: • the system • the surroundings

  7. 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.

  8. 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

  9. 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

  10. 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

  11. 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.

  12. Heat Capacity and Specific Heat • Joule-- the SI unit of heat and energy • 4.184 J = 1 cal • Specific Heat Capacity - the amount of heat it takes to raise the temperature of 1 gram of the substance by 1 oC (abbreviated “C”)

  13. 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

  14. 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

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

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

  17. Calorimetry • Calorimetry - the accurate and precise measurement of heat change for chemical and physical processes.

  18. Calorimetry • For systems at constant pressure, the heat content is the same as a property called Enthalpy (H) of the system

  19. 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)

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

  21. 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.

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

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

  24. 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

  25. 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

  26. 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

  27. 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?

  28. Summary, so far...

  29. 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

  30. 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)

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

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

  33. 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

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

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

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

  37. 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

  38. Heats of Vaporization and Condensation • Molar Heat of Vaporization(Hvap) - the amount of heat necessary to vaporize one mole of a given liquid. • Table 11.5, page 308

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

  40. 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!

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

  42. 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

  43. Section 11.4Calculating Heat Changes • OBJECTIVES: • Apply Hess’s law of heat summation to find heat changes for chemical and physical processes.

  44. Section 11.4Calculating Heat Changes • OBJECTIVES: • Calculate heat changes using standard heats of formation.

  45. Hess’s Law • If you add two or more thermochemical equations to give a final equation, then you can also add the heats of reaction to give the final heat of reaction. Called Hess’s law of heat summation • Example shown on page 314 for graphite and diamonds

  46. Why Does It Work? • If you turn an equation around, you change the sign: • If H2(g)+ 1/2 O2(g)® H2O(g) DH=-285.5 kJ • then, H2O(g) ® H2(g)+ 1/2 O2(g) DH =+285.5 kJ • also, • If you multiply the equation by a number, you multiply the heat by that number: • 2 H2O(g) ® 2 H2(g)+ O2(g) DH =+571.0 kJ

  47. Standard Heats of Formation • The DH for a reaction that produces 1 mol of a compound from its elements at standard conditions • Standard conditions: 25°C and 1 atm. • Symbol is • The standard heat of formation of an element = 0 • This includes the diatomics

  48. What good are they? • Table 11.6, page 316 has standard heats of formation • The heat of a reaction can be calculated by: • subtracting the heats of formation of the reactants from the products DHo = ( Products) - ( Reactants)

  49. CH4 (g) = - 74.86 kJ/mol O2(g) = 0 kJ/mol CO2(g) = - 393.5 kJ/mol H2O(g) = - 241.8 kJ/mol Examples • CH4(g) +2O2(g)® CO2(g) + 2 H2O(g) • DH= [-393.5 + 2(-241.8)] - [-74.68 +2 (0)] DH= - 802.4 kJ

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