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Heat Changes in Chemical Reactions

Explore the measurement and expression of heat changes in chemical reactions using calorimetry and thermochemical equations. Learn how enthalpy relates to energy changes and reactions.

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Heat Changes in Chemical Reactions

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  1. Section 11.2Measuring and Expressing Heat Changes • OBJECTIVES: • Construct equations that show theheat changes for chemical and physical processes. 1

  2. Section 11.2Measuring and Expressing Heat Changes • OBJECTIVES: • Calculate heat changes in chemical and physical processes. 2

  3. 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

  4. 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

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

  6. 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

  7. C + O2 C O2 C + O2→ CO2 + 395 kJ Energy 395kJ Reactants → Products 7

  8. O O O O C In terms of bonds C C O O 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. 8

  9. Exothermic • The products are lower in energy than the reactants • Releases energy 9

  10. CaO + CO2 CaCO3 CaCO3 + 176 kJ → CaO + CO2 CaCO3→ CaO + CO2 Energy 176 kJ Reactants → Products 10

  11. Endothermic • The products are higher in energy than the reactants • Absorbs energy • Note Figure 11.11, page 303 11

  12. 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 12

  13. 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 13

  14. 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 14

  15. 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? 15

  16. Examples • When 2 mol of solid magnesium (Mg) combines with 1 mol of oxygen gas (O2), 2 mol of solid magnesium oxide (MgO) is formed and 1204 kJ of heat is released. Write the thermochemical equation for this reaction. 16

  17. Summary, so far...

  18. 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 18

  19. Enthalpy • Symbol is H • Change in enthalpy is ΔH (delta H) • If heat is released, the heat content of the products is lower ΔH is negative (exothermic) • If heat is absorbed, the heat content of the products is higher ΔH is positive (endothermic) 19

  20. Energy Change is down ΔH is <0 Reactants → Products 20

  21. Energy Change is up ΔH is > 0 Reactants → Products 21

  22. Heat of Reaction • The heat that is released or absorbed in a chemical reaction • Equivalent to ΔH • C + O2(g) → CO2(g) + 393.5 kJ • C + O2(g) → CO2(g) ΔH = -393.5 kJ • In thermochemical equation, it is important to indicate the physical state • H2(g) + 1/2O2 (g)→ H2O(g) ΔH = -241.8 kJ • H2(g) + 1/2O2 (g)→ H2O(l) ΔH = -285.8 kJ 22

  23. Heat of Combustion • The heat from the reaction that completely burns 1 mole of a substance • Note Table 11.4, page 305 23

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