Today…

1. Today… • Turn in: • Nothing • Our Plan: • Test Results • Notes - Thermochem • Homework (Write in Planner): • Problems from the textbook (due 10/30) • You can do up through #38

2. Test Results • 11% Curve • 73.26% Average (5% higher than 1st test) • 61.1% of students scored better on this test than the first one.

3. Chemical Thermodynamics • The study of energetics of chemical reactions • Thermodynamics is the branch of science which studies the transformation of energy from one form to another

4. Why Thermodynamics? • Because everything is energy!

5. Chemical Thermodynamics • The questions that chemical thermodynamics ask are: • How much heat is evolved during a chemical reaction ("thermo chemistry")? • What determines the direction of spontaneous chemical change?

6. Example • For example, if you put a flame to a 2:1 mixture of H2 and O2 there is an almighty bang and the reaction proceeds spontaneously to products, (water) without any further supply of external energy. • Why does it not happen that water spontaneously goes to H2 and O2 when we put a lit taper to it under the same conditions?

7. Energy • Energy is the capacity to do work (to displace or move matter). • Energy literally means “work within”; however, an object does not contain work.

8. Kinetic Energy • Kinetic energy is the energy of motion. Ek = ½ mv2 m = mass v = velocity Energy has the units of joules (J or kg . m2/s2) or calories

9. Potential Energy • The energy STORED that is released through motion • It is the energy within a system because of its position or arrangement • Examples: • Ball in the air • Battery • Chemical energy

10. Potential Energy and Kinetic Energy At what point in each bounce is the potential energy of the ball at a maximum?

11. Energy Exchange • Potential energy can be converted to kinetic energy • Kinetic energy can be converted to potential energy

12. Types of Energy • Light Energy – light bulb • Heat Energy – fire • Electrical Energy – putting your finger in an electrical socket • Mechanical Energy – engine or a water wheel • Chemical Energy – burning gasoline (chemical energy is converted to heat and light)

13. Chemical Energy • Chemical energy is stored inside a substance until a substance undergoes a chemical change. Then it can be released or absorbed. • Chemical energy is stored in chemical bonds between atoms.

14. Some Important Terms • Thermodynamics – the study of heat, energy, and the ability to do work. • Energy – the capacity to do work or produce heat. • Both heat and work are energy transfers

15. Some Important Terms • The First Law of Thermodynamics – energy is not created or destroyed it is just transferred • All energy in the universe is constant!

16. Some Important Terms • Work – a force acts on something causing it to move (w) • Heat – energy transfer from thermal interaction (q)

17. Some Important Terms • System – the part of the universe we are interested in studying (we choose the system) • Surroundings – the rest of the universe

18. Types of Systems • Open: energy and matter can be exchanged with the surroundings. • Closed: energy can be exchanged with the surroundings, matter cannot. • Isolated: neither energy nor matter can be exchanged with the surroundings. After the lid of the jar is unscrewed, which kind of system is it? A closed system; energy (not matter) can be exchanged.

19. Internal Energy, E • We are going to be talking about our chemical system losing and gaining energy, so it is logical to introduce a quantity which represents the total energy of the system. This is known as the internal energy, E. • It’s important to note that your textbook refers to internal energy as U sometimes.

20. Internal Energy, E • Internal energy is defined as the total energy of our system - chemical, nuclear, heat, gravitational, any other type of energy you can think of - the sum of all the system's energy(kinetic + potential).

21. Internal Energy • A little thought will convince ourselves that it would be impossible to actually measure the total internal energy of our system. • So why define a quantity which we cannot measure? • Well, although we cannot measure the absolute value of the internal energy of a system, we can measure changes in the internal energy. And since thermodynamics is all about changes in energy this makes the change in internal energy of a system a very useful experimental quantity.

22. Changes in Internal Energy (ΔE) • The internal energy, E, of a system may change in 3 different ways. The internal energy, E, of a system will change if : • heat passes into or out of the system; • work is done on or by the system; • mass enters or leaves the system.

23. ΔE is a state function • The value of a state function is independent of the history of the system. All that matters are the starting and ending values, not how you got there. • Elevation change when climbing a mountain is an example

24. State Functions (cont.) • The fact that internal energy is a state function is extremely useful because we can measure the energy change in the system by knowing the initial energy and the final energy. • In other words, we don’t need all of the detail of a process to measure the change in the value of a state function. • In contrast, we do need all of the details to measure the HEAT or the WORK of a system.

25. Internal Energy (E) • Internal energy (E) is the total energy contained within a system • Part of E is kinetic energy (from molecular motion) • Translational motion, rotational motion, vibrational motion. • Collectively, these are sometimes called thermal energy • Part of E is potential energy • Intermolecular and intramolecular forces of attraction, locations of atoms and of bonds. • Collectively these are sometimes called chemical energy

26. Heat & Work • There are two ways to classify conversion of potential energy to kinetic energy and vice versa: • Heat – q • Work – w • Heat and work are NOT state functions (they are pathway dependent) • This makes it difficult to distinguish between the energy from heat and the energy from work. That’s again why ΔE being a state function is so important. • Car going down mountain example from Crash Course.

27. Sign Convention • Energy entering a system carries a positive sign (gains energy) • Energy leaving a system carries a negative sign (loses energy) • Energy, work, and heat are not really positive or negative. The sign just tells you the direction of flow.

28. Work (w) • A negative quantity of work signifies that the system loses energy. • - w = work done by system (expansion) • A positive quantity of work signifies that the system gains energy. • +w = work done on system (compression) • In gases, work is related to changes in volume.

29. Pressure-Volume Work For now we will consider only pressure-volume work. work (w) = –PDV

30. Heat (q) • Technically speaking, heat is not“energy.” • Heat is energy transfer between a system and its surroundings, caused by a temperature difference. Passes from higher to lower areas of heat. More energetic molecules … … transfer energy to less energetic molecules.

31. Sign Convention for Heat • + q = heat gained by system (heat is transferred to the system) • Endothermic = the system gets colder • - q = heat lost by system (heat is transferred from the system to the surroundings) • Exothermic = the system gets warmer

32. Conceptualizing an Exothermic Reaction Surroundings are at 25 °C Typical situation: some heat is released to the surroundings, some heat is absorbed by the solution. 35.4 °C 32.2 °C 25 °C In an isolated system, all heat is absorbed by the solution. Maximum temperature rise. Hypothetical situation: all heat is instantly released to the surroundings. Heat = qrxn

33. Let’s Summarize…

34. Heat • calorie - the quantity of heat required to change the temp of one gram of water one degree Celsius • SI Unit of basic energy is the Joule (J) 1 cal = 4.184 J

35. First Law of Thermodynamics • “Energy cannot be created or destroyed.” • A change in the internal energy of a system is due only to heat and work. • Inference: the internal energy change of a system is simply the difference between its final and initial states: DE = Efinal – Einitial • Additional inference: if energy change occurs only as heat (q) and/or work (w), then: DE= q + w

36. Example 6.1 A gas does 135 J of work while expanding, and at the same time it absorbs 156 J of heat. What is the change in internal energy?

37. Enthalpy (H) and Enthalpy Change (∆H) The evolved H2 pushes back the atmosphere; work is done at constant pressure. Enthalpy is the sum of the internal energy and the pressure-volume product of a system: H = E + PV Mg + 2 HCl --> MgCl2 + H2 For a process carried out at constant pressure, q=DE + PDV so q=DH Most reactions occur at constant pressure, so for most reactions, the heat evolved equals the enthalpy change.

38. Properties of Enthalpy • Enthalpy is an extensive property. • It depends on how much of the substance is present. • This makes sense because heat transfers into and out of bonds, the more moles you have, the more bonds you have, therefore more energy can be released to or taken from the environment. • Since E, P, and V are all state functions, enthalpy (H) must be a state function also. • Enthalpy changeshave unique values. DH = q Two logs on a fire give off twice as much heat as does one log.

39. Standard Enthalpy of Formation • Molar enthalpy is the amount of heat lost or gained when one mole of a compound is formed from its constituent elements. • ∆Hᵒf = Standard enthalpies of formation (molar enthalpy at 25 degrees C and 1 atm) have been calculated for nearly all compounds and are in the back of your book. (Appendix C)

40. Standard Enthalpy of Formation Examples • H2(g) + ½O2(g) → H2O(l) ΔH = –285.8 kJ mol–1 • 2C(s) + H2(g) → C2H2(g) ΔH = 227.0 kJ mol–1

41. Enthalpy Diagrams • Values of DH are measured experimentally. • Negative values indicate exothermic reactions. • Positive values indicate endothermic reactions. An increase in enthalpy during the reaction; DH is positive. A decrease in enthalpy during the reaction; DH is negative.

42. Same magnitude; different signs. Reversing a Reaction • DH changes sign when a process is reversed. • Therefore, a cyclic process has the value DH = 0.

43. Example 6.3 Given the equation (a) H2(g) + I2(s) --> 2 HI(g) DH = +52.96 kJ calculate DH for the reaction (b) HI(g) --> ½ H2(g) + ½ I2(s).

44. Example 6.4 The complete combustion of liquid octane, C8H18, to produce gaseous carbon dioxide and liquid water at 25 °C and at a constant pressure gives off 47.9 kJ of heat per gram of octane. Write a chemical equation to represent this information.

45. Enthalpy Calculations • What quantity of heat is associated with the complete combustion of 1.00 kg of sucrose (C12H22O11)? 342.3 g/mol sucrose KNOWN: -5.65 x103 kJ/mole 1 mol 342.3 g -5.65 x 103 kJ 1 mol 1000 g 1 kg x = 1.00 kg x x -1.65 x 104 kJ

46. ΔH in Stoichiometric Calculations • For problem-solving, heat evolved (exothermic reaction) can be thought of as a product. Heat absorbed (endothermic reaction) can be thought of as a reactant. • We can generate conversion factors involving DH. • For example, the reaction: H2(g) + Cl2(g) --> 2 HCl(g) DH = –184.6 kJ can be used to write: –184.6 kJ ———— 1 mol Cl2 –184.6 kJ ———— 2 mol HCl –184.6 kJ ———— 1 mol H2

47. Example 6.5 What is the enthalpy change associated with the formation of 5.67 mol HCl(g) in this reaction? H2(g) + Cl2(g) --> 2 HCl(g) DH = –184.6 kJ

48. Try it Out! • 6.5 A – What is the enthalpy change when 12.8 g H2 (g) reacts with excess Cl2 (g) to form HCl (g) H2 (g) + Cl2 (g) → 2 HCl (g) ∆H = -184.6 kJ

49. Today… • Turn in: • Get out Notes • Our Plan: • Clicker Review • Notes • Calorimetry POGIL (as practice) • Find Someone Who • Homework (Write in Planner): • Problems from the textbook (due 10/30) • You can complete up through #60

50. Heat Definitions • Heat capacity – quantity of heat required to change the temperature of a system by one degree • Molar heat capacity – quantity of heat required to change the temperature of a mole of a substance one degree