1. Terms Energy Heat Calorie Joule Specific heat Calorimeter Thermochemistry

2. Chapter 16 Thermochemistry

3. Heat (Enthalpy) Change, Δ H The amount of heat energy released or absorbed during a process.Energy Energy is the capacity to do work Energy is the capacity to do work, and can take many forms Potential energy - stored energy or energy of position Kinetic energy - energy of motion Thermal energy - (heat) movement at the atomic level

4. Law of conservation of energy Energy can be converted from one form to another, but neither created nor destroyed.

5. Heat vs. Temperature Heat – q, the energy that flows from hot to cold; Temperature – measure of average kinetic energy.

6. Calorimetry calorimeter Measures the amount of heat absorbed or released during a physical/chemical change usually by the change in temperature of a known quantity of water in a calorimeter.

7. Measuring heat Metric system: calorie (cal) –Heat required to raise the temperature of 1 gram of pure water 1ºC Food Calories differ from heat calories –1 Calorie = 1000 cal SI unit: joule (J) –1 J = 0.2390 cal –1 cal = 4.184 J

8. Use a conversion to convert nutritional Calories to calories. Converting Energy Units Use a conversion factor to convert calories to joules.

9. Practice Convert the following –934 Calories to calories –7.42x10 6 calories to Calories –236 Calories to joules –2.59x10 6 calories to joules –84 joules to calories –7.19x10 6 joules to Calories

10. Changes in State Melting Evaporation Freezing Condensation S u b l i m a t i o n

11. Heat of Fusion ( H f ) –Conversion from Solid Liquid Heat of Vaporization ( H v ) –Conversion from Liquid Gas Formula q = m H f or m H v H f water = 80 cal/g or 334 J/g H v water = 540 cal/g or 2260 J/g

12. Change in State Graph 1. Copy this graph 2. Identify specific heat, heat of fusion, and heat of vaporization.

13. Change in State Areas

14. Sample Heat of Fusion How much heat energy is needed to melt 5.0 kilograms of water at its melting point? Q = m H f = 5 kg x 334 J/kg = 1,670J Q = m H f = 5000 g x 80 cal/g = 400,000 cal

15. Collision Model Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy). Reactants must have proper orientation to allow the formation of new bonds.

16. Activation Energy The minimum energy required to transform reactants into the activated complex Flame, spark, high temperature, radiation are all sources of activation energy

17. Endo. Vs. Exo-thermic Rxn When the reaction releases heat it is exothermic, feels hot. - Δ H –Energy is given off When the reaction absorbs heat it is endothermic, feels cold. + Δ H –Energy is taken in

18. Exothermic Processes Processes in which energy is released as it proceeds, and surroundings become warmer Reactants  Products + energy

19. Endothermic Processes Processes in which energy is absorbed as it proceeds, and surroundings become colder Reactants + energy  Products

20. Endothermic Reaction w/Catalyst

21. Exothermic Reaction w/Catalyst

22. Exo vs. Endo Exothermic 1.Which has more energy, reactants or products? 2.If the reaction ran backwards, which would have more energy, reactants or products? Endothermic 1.Which has more energy, reactants or products? 2.If the reaction ran backwards, which would have more energy, reactants or products? 3.Looking at your two diagrams, which has larger activation energy? 4.In general does an exothermic reaction have a larger activation energy running forwards or backwards? 5.In general does an endothermic reaction have a larger activation energy running forwards or backwards?

23. Practice: State whether each of the following are exothermic or endothermic. 1. H + Cl  HCl + 432 kJ 2. 12 CO 2 + 11 H 2 O  C 12 H 22 O 11 + 12 O 2 ΔH=5638 kJ 3. ice  water 4. C + D  CDΔH=-65.8 kJ 5. E + F + 437 kJ  G + H 6. H 2 O vapor  H 2 O liquid

24. 7. 8. KOH  K + + OH - ΔH=-57.8 kJ 9. C 3 H 8 + 5 O 2  3 CO 2 + 4 H 2 O ΔH=-2221 kJ 10. Ca(OH) 2  Ca + O 2 + H 2 ΔH=+986 kJ 11. Fe 2 O 3 + 2 Al  Al 2 O 3 + 2 Fe ΔH=-852 kJ 12. 2 H 2 O  2 H 2 + O 2 ΔH=+572 kJ Practice: State whether each of the following are exothermic or endothermic.

25. Know: amount of heat required to raise the temperature of 1 gram by 1°C. Specific Heat That quantity is defined as the specific heat ( c ). Each substance has its own specific heat.

26. Heat absorbed/released by a substance depends on the specific heat, mass, and the temperature change. Calculating heat

27. q = the heat absorbed or released c = the specific heat of the substance m = the mass of the sample in grams ∆T is the change in temperature in °C. ∆T final temperature - the initial temperature: T final – T initial. Calculating Heat

28. Converting 256 cal into Jolues 987 cal into Joules 5.9 x10 4 J into cal 3.7 x10 4 J into cal

29. Practice Problems: q = m c ∆T 1. How much heat is lost when a piece of aluminum with a mass of 4110 g cools from 660 o C to 25 o C? (c for Al = 0.9025 J/g o C) 2. If 7.5 x 10 4 J of energy are released when 500 g of a metal cools from 100.0 o C to 20.0 o C, what is the specific heat of the metal? 3. If the temperature of 34.4 g of ethanol increases from 25.0ºC to 78.8 ºC how much heat has been absorbed? (c=.00244 J/g o C)

30. Specific Heat Practice Problems 4. A 4.50 g nugget of gold absorbed 276 J of heat. What was the final temperature of the gold if the initial temperature was 25.0 ºC and the specific heat is 0.129 J/g ºC 5. A 155 g sample of unknown substances was heated from 25.0 ºC to 40.0 ºC and it absorbed 5696 J of energy. What is the specific heat of the substance?

31. Create your own Graphs

32. Specific Heat SUBSTANCESPECIFIC HEAT CAPACITY (J/KG ºC) SUBSTANCESPECIFIC HEAT CAPACITY (J/KG ºC) Aluminum9.0 x 10 2 Alcohol (ethyl)2.3 x 10 2 Brass3.8 x 10 2 Alcohol (methyl)2.5 x 10 2 Copper3.9 x 10 2 Glycerine2.4 x 10 2 Glass (crown)6.7 x 10 2 Mercury1.4 x 10 2 Glass (pyrex)7.8 x 10 2 Nitrogen (liquid)1.1 x 10 2 Gold1.3 x 10 2 Water (liquid)4.2 x 10 3 Iron4.5 x 10 2 Water (ice)2.1 x 10 3 Lead1.3 x 10 2 Water (steam)2.0 x 10 3 Sand8.0 x 10 2 air1.0 x 10 3 Silver2.3 x 10 2