1. 1 Energy in Chemical Reactions EQ: How is energy calculated? How are reaction determined to be endo/exothermic?

2. 2 Potential Energy: Potential Energy is stored energy. n energy something has due to position or composition. What has more potential energy – a bowling ball on the ground, or a bowling ball three feet directly above your head?

3. 3 Chemical potential energy n Energy stored in the bonds between atoms of a substance. n When chemical potential energy is released, it is released as heat. n What has more potential energy – a barrel of water or a barrel of gasoline?

4. 4 Chemical potential energy n Burning isn’t the only chemical reaction that releases heat. n Virtually every chemical reaction either –Releases heat = Exothermic –Absorbs heat = Endothermic Exothermic: heat is exiting.Endothermic: heat is entering.

5. Exothermic process is any process that gives off heat – transfers thermal energy from the system to the surroundings. Endothermic process is any process in which heat has to be supplied to the system from the surroundings. System? Surroundings? The “system” is generally the container where the reaction is taking place.

6. Exothermic process is any process that gives off heat If energy is listed as a product, the reaction is exothermic: Endothermic process is any process that absorbs heat. If energy is listed as a reactant, the process is endothermic: 2H 2 ( g ) + O 2 ( g ) 2H 2 O ( l ) + energy energy + 2HgO ( s ) 2Hg ( l ) + O 2 ( g ) 6.2

7. Endothermic or Exothermic? (how to decide) Fires give off heat, and feel hot to your hands. They are exothermic. Cold packs feel cold to your hands, so they must be the opposite: endothermic. They feel cold because they’re absorbing heat from your hands.

8. 8 Chemical energy Is this reaction exothermic or endothermic?

9. 9 Potential energy Heat Is this reaction exothermic or endothermic?

10. 10 THERFORE….. n PRODUCTS in an EXOTHERMIC reaction have LESS ENERGY than REACTANTS

11. 11 Chemical energy What about this one?

12. 12 Potential energy Heat What about this one?

13. 13 THERFORE….. n PRODUCTS in an ENDOTHERMIC reaction have MORE ENERGY than REACTANTS

14. 14 Practice Questions Is this reaction endothermic or exothermic? What is the change in energy for the reaction?

15. 15 Practice Questions Does the graph represent an endothermic or exothermic reaction? What is the change in energy for the reaction?

16. 16 Energy of chemical reactions: SI unit for energy is the Joule. An energy measurement requires the following: 1. A unit ( Joules or calories). 2. A number (how much). 3. A sign to tell direction: negative = exothermic positive = endothermic Ex: 134 J or -248 J

17. 17 Unit Conversions Energy content of food is measured in calories or kilocalories (1000 calories). 0.2390 calories = 1 J or 1 calorie = 4.184 J Conversions can also be made to kJ and kcal Practice: (packet) Convert 4.42 calories to Joules

18. 18 Heat Energy Q = m x c x ΔT n Energy = Mass(m) x Specific Heat (c) x Change in Temperature (ΔT = Final – Initial) n Specific Heat: Heat required to raise the temperature by 1 degree Celcius. n Mass: Grams Specific Heat: J/(g ·⁰C), Specific Heat: J/(g ·⁰C), Temp: ⁰C Temp: ⁰C

19. 19 Practice Problem If the temperature of 34.4 g of ethanol increases from 25 °C to 78.8 °C, how much heat has been absorbed by the ethanol? The specific heat of ethanol is 2.44 J/(g · °C)

20. 20 Copper has a specific heat of 0.385 J/(g×°C). A piece of copper absorbs 5000 J of energy and undergoes a temperature change from 100 °C to 200 °C. What is the mass of the piece of copper?

21. 21 45 grams of an unknown substance undergoes a temperature increase of 38 °C after absorbing 4172.4 Joules. What is the specific heat of the substance?

22. 22 A 40 g sample of water absorbs 500 Joules of energy. How much did the water temperature change? The specific heat of water is 4.18 J/(g×°C).

23. 23 If 335 g of water at 65.5 °C loses 9750 J of heat, what is the final temperature of the water? Liquid water has a specific heat of 4.18 J/(g×°C).

24. 24 Another Phase Diagram…Temp & Energy The greater the material's internal energy, the higher the temperature of that material. Heat is the energy flow between objects of different temperature. Heat and temperature are NOT the same.

25. 25 Heat Energy & Phase Changes n Heat of Vaporization –Boiling: Liquid to Steam –Heat required to change 1g of a liquid at its normal evaporation point to a gas at the same temperature n Heat of Fusion –Melting: Solid to liquid –Heat required to change 1g of a solid at its normal melting point to a liquid at the same temperature

26. 26 Enthalpy n Heat lost or gained by a system during a physical or chemical change is called the enthalpy change (ΔH) or the heat of reaction. n Remember –A negative enthalpy change means that heat is lost from the system to the surroundings, making the process exothermic. –A positive enthalpy change tells us that heat is gained by the system and the process is endothermic.

27. 27 System Surroundings Energy  < 0 Exothermic

28. 28 System Surroundings Energy  > 0 Endothermic

29. 29 ΔH° rxn = Σ ΔH° f (products) - Σ ΔH° f (reactants) n Enthalpy of the reaction = (Sum of enthalpy of the products) MINUS (sum of enthalpy of the reactants) Calculation: Enthalpy

30. 30 SiH 4 (g) + 2 O 2 (g) ---> SiO 2 (s) + 2 H 2 O (l) Δ Δ H o f kJ/mol SiH 4 (g) 34.3 O 2 (g) 0 SiO 2 (s) -910.9 H 2 O (l) -285.8 Δ Calculate the Δ H o rxn Δ [(1 x -910.9) + (2 x -285.8)] – [(1 x 34.3) + (2 x 0)] Δ H o rxn [(1 x -910.9) + (2 x -285.8)] – [(1 x 34.3) + (2 x 0)] Δ -1482.5 – 34.3 Δ H o rxn = -1482.5 – 34.3 Δ Δ H o rxn = -1516.8 (exothermic b/c negative)

31. 31 Practice Problem: 3H 2 (g) + O 3 (g) ---> 3 H 2 O (g) Δ Δ H o f kJ/mol H 2 (g) 0 O 3 (g) 143.0 H 2 O (g) -242.2 Δ Calculate the Δ H o rxn

32. 32 Heat of Formation n Elements required to form 1 mole of a compound n Keep in mind diatomic molecules and states of matter n Ex. Heat of formation for FeCl 2 (s) Fe (S) + Cl 2 (g)  FeCl 2 (s)

33. 33 Practice Problem: n Heat of formation for the following –KBr (s) –MgO (s) –CaBr 2 (s)