Energy & Chemistry

1. Burning peanuts supply sufficient energy to boil a cup of water. Burning sugar (sugar reacts with KClO3, a strong oxidizing agent) Energy & Chemistry

2. Ch. 6 Energy and Chemistry • Vocabulary • Heat using specific heat and heat capacity • Enthalpy from single reaction • Single reagent • Limiting reagent • Enthalpy from multiple reactions • Hess’ Law • Enthalpy of reaction from ∆Hf°

3. Chapter 6: Energy and Chemical Change • Energy is the ability to do work(W = F×D or P×V) and/or supply heat • Kinetic energy is the energy an object has because of its motion • For an object of mass m with velocity v

4. The law of conservation of energy states that energy cannot be created or destroyed

5. Kinetic Energy: The Energy Of Motion KE=½mv2 Energy can be transferred by moving particles Collision of fast particles with slower particles causes the slow particle to speed up while the fast molecule slows this is why hot water cools in contact with cool air 6.1 An object has energy if it is capable of doing work 5

6. Potential Energy Breaking Bonds requires energy

7. Question: A child on a swing has the highest potential energy when: a: the swing is at its highest point b: the swing is at its lowest point c: the swing is halfway between the highest and lowest points

8. Law Of Conservation Of Energy Energy cannot be created or destroyed but can be transformed from one form of energy to another Also known as the first law of thermodynamics How does water falling over a waterfall demonstrate this law? 6.1 An object has energy if it is capable of doing work 8

9. 1st Law of Thermodynamics: For an isolated system the internal energy (E) is constant: Δ E = Ef - Ei = 0 Δ E = Eproduct - Ereactant = 0 We can’t measure the internal energy of anything, so we measure the changes in energy E is a state function E = work + heat Internal Energy is Conserved 6.2 Internal energy is the total energy of an object’s molecules 9

10. Heat And Temperature Are Not The Same The temperature of an object is proportional to the average kinetic energy of its particles—the higher the average kinetic energy, the higher the temperature Heatis energy (also called thermal energy) transferred between objects caused by differences in their temperatures until they reach thermal equilibrium 6.1 An object has energy if it is capable of doing work 10

11. Units • The SI unit of energy is the joule (J) • calorie (cal) • The dietary Calorie (note capital), Cal, is actually 1 kilocalorie

12. Energy Transfer • When a cold and hot object come into contact, they eventually reach thermal equilibrium (the same temperature) • Movie1/06_heat transfer • energy transferred as heat comes from the object’s internal energy

13. Energy Distribution

14. Energy Transfer • Thermal equilibrium: the same average KE for molecules in both objects

15. Question: As the temperature of a substance increases, what happens to the average kinetic energy of the particles? a: increases b: stays the same c: decreases

16. Question: Which sample of a substance has the higher most probable molecular speed? a: the warmer sample b: the cooler sample c: both samples have the same most probable speed

17. Definitions • current condition: state • Internal energy is a state function • state functions: independent from the mechanism or method by which a change occurred • object we are interested in: system • Everything else: surroundings • A boundary • System and Surroundings together: universe

19. Open, closed, isolated

20. Question: How is energy transferred as heat? a: energy flows from cooler objects to warmer ones. b: energy flows from warmer objects to cooler ones. c: molecules from the warmer object enter the cooler object.

21. Question: What type of system exchanges energy but not matter with the surroundings? a: open b: closed c: isolated

22. A cast iron skillet is moved from a hot oven to a sink full of water. Which of the following is not true? The water heats The skillet cools The heat transfer for the skillet has a (-) sign The heat transfer for the skillet is the same as the heat transfer for the water None of these are untrue 6.3 Heat can be determined by measuring temperature changes 22

23. The symbol for heat is “q” – two ways to calculate

24. Question: Which process is exothermic? a: N2 2N b: CO2 CO + O c: BF3 + NH3 BF3NH3

25. Ch. 6 Energy and Chemistry • Vocabulary • Heat using specific heat and heat capacity • Enthalpy from single reaction • Single reagent • Limiting reagent • Enthalpy from multiple reactions • Hess’ Law • Enthalpy of reaction from ∆Hf°

26. Heat Capacity (C) relates the heat (q) to an objects temperature change The heat capacity is the amount of heat needed to raise the object’s temperature by one degree Celsius and has the units J/°C James Joule 1818-1889

27. If my large ceramic coffee cup has a heat capacity of 35 J/ºC, how much heat does it take to raise the temperature by 50ºC James Joule 1818-1889

28. Specific Heat (s) is an intensive property, and is unique for each substance • q=m×Δt×s

29. heat gain/lost = q = (mass)(SHeat)(DT) If 25.0 g of Al (SH=.902 J/gK) cools from 310 oC to 37 oC, how many joules of heat energy are lost by the Al? where DT = Tfinal - Tinitial

30. Problem Solving A sample of nickel weighing 425 grams was initially at a temperature of 26.20C. It required 975 joules of heat energy to increase the temperature to 31.55 C. What is the specific heat of nickel?

31. Problem Solving • A 500.0 gram sample of water is initially at 25.0C. It absorbs 50.0 kJ of heat from its surroundings, what is the final temperature in C? specific heat of water = 4.184 J/gC

32. If my large ceramic cup (heat capacity = 35 J/ºC) is at 25ºC when 250 mL of 92ºC coffee is added to it. What will be the final temperature of my cup with coffee assuming no other heat loss/gain? James Joule 1818-1889

33. Problem Solving Determine the final temperature when 255 g of iron is heated to 100.0°C in boiling water and it is then quickly placed into 100.0 g water at 23.4°C. SHiron=0.449 J/gK q (water) = -q (iron)

34. Question: How much energy is required to heat 50 grams of water from 24ºC to 34ºC? a: 50 cal b: 500 cal c: 500 J

35. Question: How many dietary calories equals 1000 kJ? a: 240 b: 240,000 c: 4184 d: 4.184

36. Question: Adding 5 calories of heat energy to 1 gram of copper raises the temperature of the copper by 54.3 ºC. Adding 5 calories of heat energy to 1 gram of water raises the temperature of the water by 5.0 ºC. Which statement must be true? a: Copper has a larger specific heat capacity than water. b: Water has a larger specific heat capacity than copper. c: There must be some mistake in these measurements. d: Mass is being converted into energy in this process.

37. Question: 100 grams of water cool from 14 ºC to 13 ºC. What is the heat change for the water? a: 100 cal b: -100 cal c: 418 cal d: - 418 cal

38. work done by the system energy change FIRST LAW OF THERMODYNAMICS E = q + w heat energy transferred first law of thermodynamics, which says that energy cannot be created or destroyed

39. In reactions where gases are produce or consumed qv and qp can be very different

40. The heat produced by a combustion reaction is called the heat of combustion usually in a bomb calorimeter. The reaction is run at constant volume so that ΔE = qV Heats of reactions in solution are usually run in open containers at constant pressure, so that qP =ΔE + PΔV = ΔH

43. A calorimeter has metal parts (heat capacity of 850.0 J/degree) and 1020 grams of oil (specific heat 2.248 J/gC, both at 24.50C. Two metal slugs, one a 460.0 g piece of cobalt (specific heat 25.12 J/mol-degree), one a 360.0 g piece of cadmium (specific heat=25.34 J/mol-degree), were removed from an oven maintained at 240.0C and added to the calorimeter. If no heat was lost to the surroundings, what would be the final temperature of inside the calorimeter?

44. Question: Which equals qP? a: w b: ΔE c: ΔH d: none of these

45. Hindenburg, 6 May 1937 http://www.flatrock.org.nz/topics/flying/historic_craft_flies_again.htm http://video.google.com/videoplay?docid=-2380118142773657669&q=hindenburg

46. A sample of 50.00mL of 0.125M HCl at 22.36 ºC is added to a 50.00mL of 0.125M Ca(OH)2 at 22.36 ºC. The calorimeter constant was 72 J/g ºC. The temperature of the solution (s=4.184 J/g ºC, d=1.00 g/mL) climbed to 23.30 ºC. Which of the following is not true? qczl=67.9 J qsolution= 393.3J qrxn = 461.0 J qrxn = -461.0 J None of these 6.5 Heats of reaction are measured at constant volume or constant pressure 46

47. A sample of 50.00mL of 0.125M HCl at 22.36 ºC is added to a 50.00mL of 0.125M Ca(OH)2 at 22.36 ºC. The calorimeter constant was 72 J/g ºC. The temperature of the solution (s=4.184 J/g ºC, d=1.00 g/mL) climbed to 23.30 ºC. Which of the following is not true? qczl=67.9 J qsolution= 393.3J qrxn = 461.0 J qrxn = -461.0 J None of these 6.5 Heats of reaction are measured at constant volume or constant pressure 47

48. Ch. 6 Energy and Chemistry • Vocabulary • Heat using specific heat and heat capacity • Enthalpy from single reaction • Single reagent • Limiting reagent • Enthalpy from multiple reactions • Hess’ Law • Enthalpy of reaction from ∆Hf°

49. Energy is in chemical bonds Consider the decomposition of water: H2O(g) + 243 kJ H2(g) + 1/2 O2(g) Endothermic reaction — heat is a “reactant” DHRXN = + 243 kJ 2H2(g) + O2(g)  2H2O(g) + 486 kJ DHRXN = - 243 kJ

50. An enthalpy change for standard conditions is denoted • The physical states are important • The law of conservation of energy requires