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  1. 1. Heat
  2. 2. Homework <ul><li>Read 15.2 </li></ul><ul><ul><li>Assessment 16-22 </li></ul></ul><ul><li>Read p. 530 (Thermochemical Equations for changes of state) </li></ul><ul><ul><li>Assessment Question 27 </li></ul></ul>
  3. 3. Kinetic Theory of Matter (KTM) Kinetic Molecular Theory <ul><li>All matter is made of tiny particles that are in constant motion </li></ul><ul><li>The rate of motion determines </li></ul><ul><ul><li>the temperature of the substance </li></ul></ul><ul><ul><li>the state of matter of the substance </li></ul></ul>
  4. 4. Heat <ul><li>Heat is defined as the total amount of kinetic energy (so it depends on the amount of matter present) </li></ul><ul><li>When energy is transferred from one object to another, it is known as Heat </li></ul>
  5. 5. Temperature <ul><li>Temperature is defined as the average amount of kinetic energy </li></ul>
  6. 6. Energy and Phase Changes <ul><li>endothermic – a chemical reaction that requires energy be put in </li></ul><ul><li>- energy is a reactant </li></ul><ul><li>(more E needed to break existing bonds in the reactants than is released when the new bonds form in the products) </li></ul>
  7. 7. sunlight + 6CO 2 (g) + H 2 O(l) C 6 H 12 O 6 (aq) + 6O 2 (g) <ul><li>Photosynthesis is endothermic </li></ul><ul><li>Cooking is endothermic </li></ul>
  8. 8. <ul><li>exothermic - a chemical reaction that releases energy </li></ul><ul><ul><ul><ul><ul><li>- energy is a product </li></ul></ul></ul></ul></ul><ul><li>(more E is released than is required to break bonds in the initial reaction) </li></ul>
  9. 9. <ul><li>Reacting Zn and HCl is exothermic </li></ul>Zn(s) + HCl (aq) ZnCl 2 (s) + H 2 (g) <ul><li>Reacting Na and Cl is exothermic </li></ul>http://www.chemeddl.org/collections/whats_this/wt20080215/index.html
  10. 11. Heat Calculations <ul><li>Heat (q) </li></ul><ul><ul><li>Energy transferred from an object at a higher temperature to an object at a lower temperature. (heat lost = -heat gained) </li></ul></ul><ul><ul><li>q = mc  T </li></ul></ul><ul><ul><li>q=mH fus </li></ul></ul><ul><ul><li>q=mH vap </li></ul></ul>
  11. 12. Heat Calculations <ul><li>A 10.0g sample of iron at 50.4 o C is cooled to 25.0 o C in 50.0g of water. Calculate the amount of heat lost by the iron. </li></ul><ul><li>c iron= 0.449 J/g o C </li></ul><ul><li>A 2.1g ice cube at –8.0 o C melts completely and warms to 12.5 o C. How much heat was required? </li></ul><ul><li>H fus ice = 334 J/g </li></ul><ul><li>c ice = 2.03 J/g o C </li></ul><ul><li>c water = 4.18J/g o C </li></ul>
  12. 13. Some Helpful Constants <ul><li>specific heats: </li></ul><ul><li>gold = 0.129 J/g ° C </li></ul><ul><li>aluminum = 0.897 J/g ° C </li></ul><ul><li>water = 4.184 J/g ° C </li></ul><ul><li>iron = 0.449 J/g ° C </li></ul><ul><li>silver = 0.135 J/g ° C </li></ul><ul><li>ice = 2.03 J/g ° C </li></ul><ul><li>Hfus water = 334 J/g </li></ul><ul><li>Hvap water = 2260 J/g </li></ul>
  13. 14. Practice <ul><li>How much heat is required to raise the temperature of 68.0g of aluminum from room temperature 24.0 ° C to 80.0 ° C? </li></ul><ul><li>A beaker containing 150.0g of water at 80.0 ° C is cooled to 50.0 ° C. How much heat is lost by the water? </li></ul><ul><li>A 75.0g chunk of ice at –15.0 ° C is warmed to 0.0 ° C, completely melted, then warmed to 25.0 ° C. How much energy did this process require? </li></ul><ul><li>If a 1.55g piece of stainless steel at 30.0 ° C absorbs 141.0 J of heat, what will its final temperature be? </li></ul><ul><li>An unknown sample of metal has a mass of 0.625g. It is heated from 25.0 ° C to 55.0 ° C with 2.42 J of energy. What is the specific heat of this metal? Identify the metal. </li></ul>