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Thermochemistry

This review covers thermochemistry topics such as temperature, heat, energy changes in chemical reactions, and phase changes including melting, boiling, condensation, freezing, and vaporization. It also includes examples and calculations related to specific heat, latent heats, and heating/cooling curves.

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Thermochemistry

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  1. Thermochemistry

  2. Phase Change Review • Melting: solid  liquid • Vaporization: liquid  gas • Evaporation: occurs only at surface of liquid • Boiling: bubbles of vapor form within the liquid when the temperature of the liquid is increased • Condensation: gas  liquid • Freezing: liquid  solid

  3. Thermochemistry • Thermochemistry: study of energy changes that occur during chemical reactions and changes in physical state.

  4. How does Heat differ from Temperature? • Temperature: average kinetic energy of particles • The faster the particles are moving, the higher the temperature. • Heat (q): transfer of kinetic energy from an object at a higher temperature to an object at a lower temperature. • Units: Joules or calories (kCal = Cal in foods) • 1 calorie = 4.184 J

  5. Joule/Calorie Conversion Practice • Example: 9.6 calories = ______ Joules • 4500 Joules = ______ calories • 53 calories = ______ Joules • 2310 Joules = _______ Calories *Challenge!!!

  6. Energy Changes in Chemical Reactions • Chemical Potential Energy (PE) – Energy stored in chemical bonds. • Exothermic Reactions • Endothermic Reactions

  7. Exothermic Reactions • Exothermic Reactions: net release of energy (heat) • “feels hot to the touch” • ex. combustion reactions • PE of Reactants > PE of Products Reactants Activation Energy Potential Energy (J) Products Enthalpy Time

  8. Endothermic Reactions • Endothermic Reactions: net absorption of energy (heat) • “Feels cold to the touch” • ex. decomposition of water using electrolysis • PE of Reactants < PE of Products Reactants Activation Energy Potential Energy (J) Products Enthalpy Time

  9. Enthalpy • Enthalpy (ΔH): the difference in energy between the products and reactants in a chemical change • ΔHreaction = Hproducts - Hreactants • ΔHreactionis negative for exothermic reactions. Why? • ΔHreactionis positive for endothermic reactions. Why? Example: • 2H2 (g) + O2 (g)  2H2O (l) ΔH = -572 kJ • 2H2O (l)  2H2 (g) + O2 (g) ΔH = +572 kJ

  10. Law of Conservation of Energy • Law of Conservation of Energy: the energy released when elements come together to form a compound is equal to the energy required to decompose the compound into its elements. Example: • 2H2 (g) + O2 (g)  2H2O (l) ΔH = -572 kJ • 2H2O (l)  2H2 (g) + O2 (g) ΔH = +572 kJ

  11. Activation Energy & Catalysts • Activation energy: The amount of energy needed to get a reaction started. • Required for BOTH endothermic and exothermic reactions.

  12. How does a catalyst speed up a reaction? • Catalyst: speeds up the reaction, but is not used up during the reaction. Activated Complex Activation Energy Catalyst Reduces Activation Energy Reactants Potential Energy (J) Products Time

  13. Potential Energy Diagram WS

  14. Endothermic or Exothermic? • Reactants  Products + Energy • Reactants + Energy  Products • 2NaHCO3 (s) + 129 kJ  Na2CO3 (s) + H2O (g) + CO2 (g) • CaO (s) + H2O (l)  Ca(OH)2 (s) + 65.2 kJ • baking bread • campfire • burning fuel • photosynthesis • breaking down carbohydrates for energy • Electrolysis of water

  15. Endothermic or Exothermic? • Phase changes are physical changes (no new substance is formed)… • Melting • Freezing • Boiling/Evaporation/Vaporization • Condensation

  16. Thought Question • Why are orange trees sprayed with water before cold nights to protect the fruit from frost damage?

  17. Specific Heat • Specific Heat (C): amount of energy (in Joules) needed to increase the temperature of 1 gram of a substance by 1oC. • See Reference Tables q = mCΔT • q = heat (joules or calories) • m = mass (grams) • C = specific heat (J/g*oC) • ΔT = change in Temperature = Tf - To

  18. Remember!!! • Temperature does not change during a phase change! (Energy goes into pulling molecules farther apart.) • Therefore, we cannot use this equation for phase changes: q = mCΔT

  19. Latent Heats • Heat of Vaporization (Hv): amount of energy absorbed when one gram of a liquid is changed into a vapor (or, the amount of energy released when 1 gram of a vapor is changed into a liquid). q = mHv • Hv = heat of vaporization (J/g) • Hv of water = 2260 J/g

  20. Heat of Fusion • Heat of Fusion (Hf): amount of energy needed to convert 1 gram of a solid into a liquid. q = mHf • Hf = heat of fusion (J/g) • Hf of water = 334 J/g

  21. Heating and Cooling Curves • Label Diagram • +q  heat absorbed • -q  heat released

  22. Example Problems – Level One • How much heat is added if 100 grams of liquid water increases in temperature from 30oC to 70oC? • How much heat is absorbed if 200 g of ice increases in temperature from -15oC to -5oC?

  23. (cont’d) • How much heat is released if 80 grams of water vapor decreases in temperature from 150oC to 125oC? • How much heat is absorbed when 30 g of ice is changed into liquid water at 0oC? • How much heat is released when 50 grams of water vapor is changed into liquid water at 100oC?

  24. Example Problems – Level Two • How much heat is absorbed if 30 grams of water at -10oC is converted into liquid water? • How much heat is absorbed if 45 grams of water at 80oC is converted into steam at 105oC?

  25. (cont’d) • How much heat is released if 500 grams of vapor at 120oC changes to liquid water 70oC?

  26. Calorimetry • Calorimetry: the precise measurement of the heat flow out of a system for chemical and physical processes. • Calorimeter: composed of a reaction chamber surrounded by water. • The water supplies heat for an endothermic reaction or absorbs heat from an exothermic reaction. • Use q = mCΔT to find the amount of heat released or absorbed by the reaction.

  27. Calorimeter

  28. Calorimeter Ex Problems (short) NOT in packet • How much heat is released by the reaction if 10 grams of water is heated from 20oC to 60oC. • What is the final temperature of 15 g of water if the water starts at 20oC and loses 300 J of heat? • Determine the substance if a sample of 200 grams increases in temperature from 10oC to 28oC when it absorbs 1620 Joules of energy.

  29. Calorimeter Example Problems NOT in packet • How much heat is released by the reaction if 10 grams of water is heated from 20oC to 60oC. • What is the mass of the water if a release of 650 J of heat causes the water to increase in temperature by 15oC?

  30. Example Problems • What is the final temperature of 15 g of water if the water starts at 20oC and loses 300 J of heat? • Determine the substance if a sample of 200 grams increases in temperature from 10oC to 28oC when it absorbs 1620 Joules of energy.

  31. Entropy • Entropy (S): a measure of the disorder or randomness of the particles that make up a system. Observations: • Systems tend to go from a state of high energy to a state of low energy (exothermic reactions are more likely to occur than endothermic reactions). • Systems tend to become more disordered. • Ex. a deck of cards, tea cup

  32. Spontaneous Reactions • A reaction is spontaneous if: • Energy decreases (energy is released = exothermic)AND • Disorder increases (positive entropy). Ex. Wood burning

  33. Entropy is NOT conserved • Entropy is NOT conserved. The natural tendency of entropy is to increase. In most spontaneous, naturally occurring processes, entropy increases. • ΔSsystem = Sproducts - Sreactants

  34. Predicting Changes in Entropy • Predict whether ΔSsystem will be positive or negative: 1. H2O (l)  H2O (g) [positive] 2. CH3OH (s)  CH3OH (l) [positive] 3. Increase the temperature of a substance [positive]

  35. (cont’d) 4. Dissolving of a gas in a solventCO2 (g)  CO2 (aq) [negative] 5. What happens when the number of gaseous product particles is greater than the number of gaseous reactant particles?2SO3 (g)  2SO2 (g) + O2 (g) [positive] 6. Generally, when a solid or liquid dissolves to form a solution?NaCl (s)  Na+ (aq) + Cl- (aq) [positive]

  36. (add….) • MP / BP chart… when is it a solid, liquid, or gas?

  37. Bell Work for Monday, May 14 • Grab a scantron, periodic table, Bohr model sheet • Calculator (if needed) • Scantron • Name • Subject: Thermochemistry Test • Date • Study for your test!!

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