Thermodynamics

1. Thermodynamics Honors Unit 5

2. Energy: Basic Principles • Thermodynamics – the study of energy changes • Energy – the ability to do work or produce heat • Note: Work is force acting over a distance

3. Energy: Basic Principles • Kinetic Energy – energy of motion • KE = • Potential Energy – energy due to position or composition

4. Law of Conservation of Energy • A.k.a. first Law of Thermodynamics • Energy can be converted from one form to another but can’t be created or destroyed • This means the total energy of the universe is CONSTANT!

5. Heat vs. Temperature • Temperature – measure of the random motion of a substance • Temperature is proportional to kinetic energy (it is a measure of the average kinetic energy in a substance) • Heat (q) – flow of energy due to a temperature difference

6. Important Aspects of Thermal Energy & Temperature • Heat is NOT the same as temperature! • The more kinetic energy a substance has, the greater the temperature of its atoms and molecules. • The total thermal energy in an object is the sum of its individual energies of all the molecules. • For any given substance, its thermal energy depends not only on its composition but also on the amount of substance

7. System vs. Surroundings • A system is the part of the universe we are studying. • The surroundings are everything else outside of the system.

8. Direction of Heat Flow • Heat transfer occurs when two objects are at two different temperatures. • Eventually the two objects reach the same temperature • At this point, we say that the system has reached equilibrium.

9. Thermal Equilibrium • Heat transfer always occurs with heat flowing from the HOT object to the COLD object.

10. Thermal Equilibrium • Transfer of heat continues until both objects are at the same temperature!

11. Thermal Equilibrium The quantity of heat lost by the hotter object and the quantity of heat gained by the cooler object are EQUAL.

12. Exothermic vs. Endothermic • Exothermic process heat is transferred from system to the surroundings • Heat is lost from the system (temperature in system decreases) • Endothermic process  (opposite of exothermic process) heat transferred from surroundings to the system • Heat is added to the system (temperature in system increases)

13. Exothermic Process

14. Endothermic Process

15. Units of Energy • Joule (J) is the SI unit of energy & heat One kilojoule (kJ) = 1000 joules (J) • calorie (cal) = heat required to raise the temperature of 1.00 g of water by 1 °C 1 calorie = 4.184 J

16. Units of Energy • Food is measured in Calories (also known as kilocalories) instead of calories 1 Cal = 1 kcal = 1000 calories

17. Units of Energy 3800 cal = __________ Cal = _________ J

18. Units of Energy The label on a cereal box indicates that 1 serving provides 250 Cal. What is the energy in kJ?

19. Heat Transfer Direction and sign of heat flow – MEMORIZE! ENDOTHERMIC: heat is added to the system & the temperature increases (+q) EXOTHERMIC: heat is lost from the system (added to the surroundings) & the temperature in the system decreases (-q)

20. Heat Capacity • The quantity of heat required to raise an object’s temperature by 1 °C (or by 1 Kelvin) • Heat capacity is an extensive property. • Which will take more heat to raise the temperature by 1 °C?

21. Specific Heat (Specific Heat Capacity) • Specific Heat (C) - The quantity of heat required to raise the temperature of one gram of a substance by 1 °C • Intensive property • Units: • J/(g°C) or J/(g°K) • cal/(g°C) or cal/(g°K)

22. Examples of Specific Heat At the beach, which gets hotter, the sand or the water? Higher specific heat means the substance takes longer to heat up & cool down!

23. Examples of Specific Heat • Specific heat (C)= the heat required to raise the temperature of 1 gram of a substance by 1 °C Cwater = 4.184 J/(g°C) Csand = 0.664 J/(g°C)

24. Calculating Changes in Thermal E q = m x C xΔT q = mCΔT q = heat (cal or J) m = mass (g) C = specific heat capacity, J/(g°C) ΔT = change in temperature, Tfinal – Tinitial (°C or K) ***All units must match up!!!***

25. Example q = mCΔt How much heat in J is given off by a 75.0 g sample of pure aluminum when it cools from 84.0°C to 46.7°C? The specific heat of aluminum is 0.899 J/(g°C).

26. Example q = mCΔt What is the specific heat of benzene if 3450 J of heat are added to a 150.0 g sample of benzene and its temperature increases from 22.5 °C to 35.8 °C?

27. Example q = mCΔt A 50.0 g sample of water gives off 1.025 kJ as it is cooled. If the initial temperature of the water was 85.0 °C, what was the final temperature of the water? The specific heat of water is 4.18 J/(g°C).

28. Calorimetry • Calorimetry: measurement of quantities of heat • A calorimeter is the device in which heat is measured.

29. Calorimetry • Assumptions: • Heat lost = -heat gained by the system • In a simple calorimeter, no heat is lost to the surroundings

30. Coffee Cup Calorimetry ***Use styrofoam instead of a beaker to keep heat in Steps: • Add hot solid metal to cool water • Water will heat up (T rises) as metal cools • Eventually, water & metal are at same T. qmetal + qwater = 0 qmetal = -qwater (Heat lost by metal = -heat gained by the water)

31. Calorimetry qmetal = -qwater Heat lost by metal = -heat gained by water Since q = mCΔT, mmCmΔTm = -[mH2OCH2OΔTH2O]

32. Sample Problem A 358.11 g piece of lead was heated in water to 94.1 °C. It was removed from the water and placed into 100. mL of water in a Styrofoam cup. The initial temperature of the water was 18.7 °C and the final temperature of the lead and water was 26.1 °C. What is the specific heat of lead according to this data?

33. A 358.11 g piece of lead was heated in water to 94.1 °C. It was removed from the water and placed into 100. mL of water in a Styrofoam cup. The initial temperature of the water was 18.7 °C and the final temperature of the lead and water was 26.1 °C. What is the specific heat of lead according to this data?

34. Bomb Calorimeter • Constant volume “bomb” calorimeter • Burn sample in O2 • Some heat from reaction warms water • qwater = mCH2OΔT • Some heat from reaction warms bomb • qbomb = CbombΔT • qrxn + qH2O + qbomb = 0

35. Energy & Changes of State • All changes of state involve energy changes (more in Unit 9) • Note that fusion = melting

36. State Functions • A property where the change from initial to final state does not depend on the path taken • Ex.) The change in elevation from the top to bottom of a ski slope is independent of the path taken to go down from the slope

37. Enthalpy Changes for Chemical Rxns. • Heat of reaction • The heat absorbed or given off when a chemical reaction occurs at constant T (temp.) and P (pressure)

38. Enthalpy • Enthalpy (H) • The heat content of a reaction (chemical energy) • ΔH = change in enthalpy • The amount of energy absorbed by or lost from a system as heat during a chemical process at constant P • ΔH = Hfinal - Hinitial

39. Properties of Enthalpy • Enthalpy is an extensive property • It does depend on quantity • Enthalpy is a state function • Depends only on the final & initial values • Every reaction has a unique enthalpy value since ΔH = Hproducts - Hreactants

40. Representation of Enthalpy as a Graph

41. Two Ways to Designate Thermochemical Equations Endothermic: • H2 (g) + I2 (s)  2 HI (g) ΔH = 53.0 kJ • H2 (g) + I2 (s) + 53.0 kJ  2 HI (g)

42. Two Ways to Designate Thermochemical Equations Exothermic: • ½ CH4 (g) + O2 (g)  ½ CO2 (g) + H2O (l)ΔH = -445.2 kJ • ½ CH4(g) + O2(g)  ½ CO2(g) + H2O(l) + 445 .2 kJ

43. Two Ways to Designate Thermochemical Equations Note the meaning of the sign in ΔH in the equations above!! Endothermic: ΔH = + Exothermic: ΔH = -

44. Two Ways to Designate Thermochemical Equations Note the important of designating the physical state or phase of matter. Why?? Because this will change the heat of reaction! (ΔH)

45. Thermochemical Equations What do the coefficients stand for? How can they differ from the ones we have used before? Coefficients = the number of moles (as before) BUT We can use fractional coefficients now!

46. Thermochemical Equations What is the standard state? How do we designate conditions of temperature and pressure that are not at standard state? Standard state = 1 atm pressure & 25 °C ΔH° = ΔH at standard state Must show conditions over arrow if not at standard state!

47. Thermochemical Equations How can we find the enthalpy of reaction when we reverse it? Reverse the reaction, reverse the sign of ΔH! Example: CO (g) + ½ O2 (g)  CO2 (g) ΔH = -283 kJ CO2 (g)  CO (g) + ½ O2 (g) ΔH = +283 kJ

48. Example Given Rxn. #1, find the ΔH for Rxns. 2 & 3 Reaction #1 2 SO2 (g) + O2 (g)  2 SO3 (g) ΔH = +197.8 kJ Reaction #2 SO2 (g) + ½ O2 (g)  SO3 (g) ΔH = Reaction #3 4 SO3 (g)  4 SO2 (g) + 2 O2(g) ΔH =

49. ΔH as a Stoichiometric Quantity Given the reaction below, how much heat is produced when 15.0 g of NO2 are produced? 2 NO (g) + O2 (g)  2 NO2 (g) ΔH = -114.1 kJ

50. ΔH as a Stoichiometric Quantity Given: ΔH = -283 kJ CO (g) + ½ O2 (g)  CO2 (g) (a) Calculate the enthalpy of the above reaction when 3.00 g of product are formed