Thermodynamics and Phase Changes

1. Thermodynamics and Phase Changes

2. Energy • Energy – E – the combination of the amount of work and heat a sample is able to transfer • The study of energy changes is thermodynamics • Energy is measured in Joules • Work – w – the energy required to move an object by applying a force over a distance • Chemists generally don’t do work. • Heat – q – the transfer of energy due to a difference in temperature • Heat is a verb and not a noun in science • Heat and Temperature are not the same thing.

3. Heat vs. Temperature • Heat and Temperature are not the same thing. • Heat is a measurement of energy flow from one object to another in Joules. • Temperature is a measurement that allows us to determine if heat CAN flow from one object to another. (Different temperatures or the same temperature) • Heat is energy flow. • Temperature is an average energy amount of an object.

4. First Law of Thermodynamics • ΔEuniverse = 0 • The amount of energy in the universe is constant. Energy is not created or destroyed.

5. Thermodynamic Conventions • A few standard definitions and concepts • Universe – everything in existence • System – the part of the universe that is being studied • Surroundings – everything in the universe except the system

6. Thermodynamic Conventions • All thermodynamic numerical values have a number and a sign. • The sign shows the direction of energy flow • A positive sign means energy is added to the system. • A negative sign means energy is released from the system.

7. Reactions • Consider the combustion of methane • Feels hot • Transfer of heat from reaction to surroundings • q = negative • Exothermic – heat “leaves” the system and goes into the surroundings

8. Reactions • Consider an instant ice pack • Feels cold • Heat is transferred from surroundings into the system. • q = positive • Endothermic – heat “goes into” the system from surroundings

9. Enthalpy • Enthalpy = H • If we consider a system at constant pressure: • Everyday existence • ΔH = Hfinal – Hinitial = q = heat

10. Heat • When no phase changes or reactions are involved Where: • q = heat (in Joules) • m= mass • Cp = specific heat – energy required to change the temperature of one gram of substance by 1°C or 1K (units = J/g °C or J/g K) • ΔT = change in temperature = Tfinal - Tinitial

11. Practice Problems • How much heat in kJ is required to heat a 100.0g sample of water from 20.0°C to 80.0°C?

12. Practice Problems • If 330J of heat is removed from a 10.0g block of zinc at 20.0°C, what will be the final temperature?

13. Determining Heat Transfer • Calorimetry – process for measuring heat transfers • Uses the temperature change of an object with a known mass, and specific heat to calculate the heat absorbed or released in a process. • Usually water is used as the known object • Sometimes referred to as coffee cup calorimetry • Under constant pressure q = ΔH

14. Practice A 4.57g sample of an unknown metal is heated in boiling water bath at 98.1°C. The metal is then placed in a coffee cup calorimeter with 15.20g of water which is initially at 22.3°C. The mixture’s temperature peaks at 27.5°C. What is the specific heat of the metal?

15. Practice A 9.31g piece of an unknown metal is placed in a boiling water bath at 99.3°C. It is then placed in a coffee cup calorimeter with 25.31g of water. The initial temperature of the water in the calorimeter is 24.1°C and it rises to 27.4°C once the metal is placed in. What is the identity of the unknown metal?

16. Phase Changes

17. Phase Changes • Relate heat to temperature change • q = mCpΔT • Applies where there is a change in temperature. • AB, CD, EF

18. Phase Changes • To change solid to a liquid • Weaken some of the intermolecular forces • Input of energy • Enthalpy of fusion, Hf • q = mHf

19. Phase Changes • To change liquid to a gas • Totally break the intermolecular forces • Input of energy • Enthalpy of vaporization, Hv • q = mHv

20. Practice Problems • How much heat is required to melt 50.0g of ice at 0.0°C?

21. Practice Problems • What heat flow is produced by condensing 25.0g of steam at 100.0°C to water?

22. Practice Problems • What is the heat flow for converting 105g of steam at 120°C to ice at -15°C?

23. Water in A Vacuum Pump

24. Phase Diagram • Triple Point – where all three states exist together. • Critical Point – point where a gas can no longer be liquefied = supercritical fluid

25. Water Phase Diagram

27. Carbon Dioxide Phase Diagram • Dry ice sublimes normally because at 1 atm you find the transition between solid and gas.

28. Phase Diagrams

29. Phase Diagrams • The slope of the liquid-solid boundary is negative in water • Liquid water is more dense than ice. • This is not normal.

30. Reaction Pathway Diagrams

31. Reaction Pathway Diagrams • Plot energy of substance versus its place in the progress of the reaction • Reaction coordinate – reaction progress

32. Reaction Pathway Diagrams • Consider Bunsen Burner flame • What term do we use to describe reactions that “feel” hot? • Exothermic – releases heat to surroundings

33. Reaction Pathway Diagrams • Endothermic – absorbs heat from surroundings • Feels cold • Ice packs = endothermic reaction

34. Activation Energy • Reactions need a certain amount of energy to start them. • Sometimes it is very small, sometimes it is very large • Activation energy (Ea) – energy needed to start a reaction

35. Transition State • Transition state – the high energy state between the products and reactants • Transition states are unstable, they always go on to form products or decay back into starting materials

36. Labeling Reaction Diagrams

37. Labeling Reaction Diagrams

38. Labeling Reaction Diagrams

39. Labeling Reaction Diagrams

40. Catalysis • Catalyst – compound that speeds up a reaction but is not consumed by the reaction

41. Catalysis • Catalysts must work by changing the path a reaction goes through. • Black is the normal reaction profile • Orange is the catalyzed profile • Catalysis changes the transition states and lowers activation energy