Energy (Ch. 10)

1. Energy(Ch. 10)

2. What is Energy? 10.1 Energy, Temperature and Heat Energy is the ability of something to do work or make heat.

3. Forms of Energy • Heat • Light • Sound • Motion • Electrical

4. Energy Change • The Law of Conservation of Energy: Energy cannot be made or destroyed, but it can change from one type to another type (of energy). • If you ride a bicycle down a hill, you move faster and faster. When the hill ends, you slow down and stop. Why? • You slow down because friction of the tires on the road changed your energy to heat. • The energy change: Motion  Heat

5. Work • When you ride your bicycle down the hill, gravity does “work” on you. • Work is a force pushing or pulling on something over distance.

6. Work + Heat = Energy • Case A: You ride down a hill on your bicycle and go very fast. • Case B:Your friend uses breaks, goes slow and stops at the bottom. • For both people, the energy change is the same! • Energy is a state function because it does not depend on the “pathway” (how the energy change happened). • Speed is NOT a state function, because it depends on

7. Temperature and Heat • Temperature is the measure of random motions of the things. • Heat is the movement of energy. • Thermal energy is the random movement of molecules and atoms.

8. Temperature and Heat • If you have a hot cup of coffee, after about an hour it will cool to be the same temperature as the room. • What happened to the molecules in the coffee and the air? • At first, the molecules in the coffee are moving faster than the molecules of air. • The coffee molecules move slower. • The air molecules move faster. • The energy moves from the coffee to the air.

9. Exothermic Vs. Endothermic • The system is the thing we are interested in. • The surroundings are everything else in the world. • Example: I want to measure the temperature change of my cup of coffee. • The system = the coffee cup • The surroundings = everything else (the air, chairs in the room, the students…)

10. Exothermic Vs. Endothermic • Exothermic = energy being released • Endothermic = energy being absorbed

11. Exothermic Vs Endothermic REACTIONS To determine if a chemical reaction is exothermic or endothermic you must look at the energy of the whole reaction • Exothermic Reaction – OVERALL, the reaction releases more energy than it takes in • Endothermic Reaction – OVERALL, the reaction absorbs more energy than it releases

12. Energy of the Molecules After the Reaction Energy of Molecules before the Reaction How this looks on a graph Energy released to the surroundings

13. Does this Graph Represent an Exothermic or Endothermic Reaction? • Exo or Endo? • Exothermic, because molecules have less energy then when they started

14. Does this Graph Represent an Exothermic or Endothermic Reaction? • Exo or Endo? • Endothermic, because molecules have more energy then when they started

15. 10.2 The Flow of Energy

16. Thermodynamics • Thermodynamics is the study of energy • The first law of thermodynamics = the law of conservation of energy: the energy of the universe is constant. • Internal energy (E) = energy of the system Δ E = q + w Δ E means the change in energy q = heat w = work

17. Measuring Energy Changes • You can measure heat in joules (J) or calories (cal) • One calorie = energy needed to raise the temperature of one gram of water one degree Celsius. 1 cal = 4.184 J Practice: Convert 44.2 cal to J

18. Thermodynamic quantities • Thermodynamic quantities (q, w and Δ E) have two parts • Number = the size of the change • Sign = (+ or – )shows the direction of the flow.

19. Calculating Heat from lab data Energy = Specific heat x Mass x Temp change Q = smΔT This equation tells us how much energy was gained by a substance (if it was heated) or lost by a substance (if it was cooled) What part of this equation will tell you if something is heated or cooled? Temp change (also written as ∆T)

20. What is a “Specific Heat”? Specific Heat Capacity – the amount of energy required to heat 1 gram of something by 1°C. Every substance has a different specific heat What is the specific heat of water… 4.18 J/(g*°C)  memorize! This number tells us that if we want to raise the temperature of a container of water by 1°C, then each gram of water in the container must gain 4.18 J of energy.

21. Let’s do an example • Let’s say your making some delicious Ramen noodles for dinner mmm..mmmm good! • The directions say to boil 2 cups of water (~500 grams of water) • If the water starts at room temp (~ 20°C), how much energy must it gain to boil?

22. 10.3 Thermochemistry • Enthalpy = ΔH = the heat for a reaction • The units of Enthalpy are kJ • Example: If you burn one mole of CH4 the enthalpy (ΔH) is 890 kJ. • In other words: CH4 + O2 CO2 +H2O ΔH = -890 kJ How much heat is released if you burn 5.8 grams methane?

23. An exothermic reaction has a negative (-) enthalpy. • Is this reaction exothermic or endothermic? CH4 + O2 CO2 +H2O ΔH = -890 kJ • Yes, exothermic • The reverse reaction would be endothermic and so it has a positive (+) sign: CO2 + H2O  CH4 + O2 ΔH = 890 kJ

24. Calorimeter • A calorimeter is a device used to measure the heat in a chemical reaction.

25. Hess’s Law • Hess’s Law: The change in enthalpy for a reaction is the same whether a reaction takes place in one step or many steps. • Example: What is the change in enthalpy for: N2 + O2 NO2 ΔH = ???

26. N2 + O2 NO2 ΔH = ??? • You know: N2 + O2 2NO ΔH = 180 kJ 2NO2  2NO + O2ΔH = 112 kJ Can we make NO2 from N2 and O2 in two steps? Yes, but…you have to do the second reaction backwards… 2NO + O2 2NO2ΔH = -112 kJ

27. If we add these two reactions, we can find the ΔH of N2 + O2 NO2 N2 + O2 2NO ΔH = 180 kJ 2NO + O2 2NO2ΔH = -112 kJ N2 + O2 NO2 ΔH = 180 + (-112) = 68 kJ

28. 10.4 Using Energy in the Real World • Read p339-349 and answer questions 1-7 on p350 • Go to this websites • http://www.adventuresinenergy.org/ - “What are Oil and Natural Gas?” and “Refining: Distillation” • http://science.howstuffworks.com/oil-refining4.htm • And answer the questions on the worksheet, “Petroleum Assignment”