1. Thermodynamics X Unit 9

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

7. 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.

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

9. Exothermic vs. Endothermic  Exothermic process  heat is transferred from the system to the surroundings Heat is lost from the system (temperature in system decreases)  Endothermic process  heat transferred from the surroundings to the system Heat is added to the system (temperature in system increases)

10. Exothermic Process Products have lower energy than reactants Energy products < Energy reactants

11. Endothermic Process Reactants have lower energy than products Energy products > Energy reactants

12. 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 = 4.18 J

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

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

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

16. 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 )

17. 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  Units: J/(g · °C) or J/(g · K) cal/(g · °C) or cal/(g · K)

18. 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!

19. Examples of Specific Heat  Specific heat (C)= the heat required to raise the temperature of 1 gram of a substance by 1 °C C water = 4.184 J/(g  °C) C sand = 0.664 J/(g  °C)

20. Calculating Changes in Thermal E q = mCΔT q = heat (J) m = mass (g) C = specific heat capacity, J/(g  °C) ΔT = change in temperature, T final – T initial (°C or K) ***All units must match up!!!***

21. Example 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). q = mCΔt

22. Example 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? q = mCΔt

23. Calorimetry  Calorimetry: measurement of quantities of heat A calorimeter is the device in which heat is measured.

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

25. 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 = ΔH final - ΔH initial

26. Representation of Enthalpy as a Graph

27. Two Ways to Designate Thermochemical Equations Endothermic: a) H 2 (g) + I 2 (s)  2 HI (g) ΔH = 53.0 kJ b) H 2 (g) + I 2 (s) + 53.0 kJ  2 HI (g)

28. Two Ways to Designate Thermochemical Equations Exothermic: a) ½ CH 4 (g) + O 2 (g)  ½ CO 2 (g) + H 2 O (l) ΔH = -445.2 kJ b) ½ CH 4 (g) + O 2 (g)  ½ CO 2 (g) + H 2 O(l) + 445.2 kJ

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