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Week of 9/20-9/22

Week of 9/20-9/22. Bell work. Turn in your homework from last class (Chapter 5 # 50, 52, 56, 62, 64, 72, 74) Calculate  H for the reaction given the following information:. Agenda. Bell work Revisit end of chapter 5 Chapter 5 problem solving (30 minutes) Chapter 6 – main points

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Week of 9/20-9/22

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  1. Week of 9/20-9/22

  2. Bell work Turn in your homework from last class (Chapter 5 # 50, 52, 56, 62, 64, 72, 74) Calculate H for the reaction given the following information:

  3. Agenda • Bell work • Revisit end of chapter 5 • Chapter 5 problem solving (30 minutes) • Chapter 6 – main points • Homework – read chapter 6 and complete overview

  4. Looking ahead… • Thursday, Sept 22 • Chapter 6 problem solving • Overview of chapter 7 • Homework – read chapter 7 and complete problems from chapters 6 and 7 (due Thursday, October 14) • Tuesday, • Test over chapters 4 and 5 • Thursday, October • Due: Chapter 7 overview and problems from chapters 6 and 7 • Tuesday, October (after fall break) • Midterm (chapters 1-7 covered) • Thursday, October • Calorimetry lab

  5. Practice Calculate H for the reaction given the following chemical equations and their respective enthalpy changes:

  6. Enthalpy changes Enthalpies of vaporization (ΔH for converting liquids to gases) Enthalpies of fusion (ΔH for melting solids) Enthalpies of combustion (ΔH for combusting a substance in oxygen) Enthalpy of formation (or heat of formation; ΔH for formation of a compound from its constituent elements)

  7. Enthalpy of formation • ΔHf • Magnitude of any enthalpy change depends on the conditions of temperature, pressure, and state of reactants and products • Because of this, we need standard conditions (25°C and 1.00 atm pressure). • Standard enthalpy change: H° • By definition, the standard enthalpy of formation of the most stable form of any element is zero because there is no formation reaction needed when the element is already in its standard state

  8. Standard enthalpies of formation

  9. Calculation of H We can use Hess’s law in this way: H = nHf(products) - mHf(reactants) where n and m are the stoichiometric coefficients.

  10. Calculation of H C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (l) • H = nHf(products) - mHf(reactants) H = [3(-393.5 kJ) + 4(-285.8 kJ)] - [1(-103.85 kJ) + 5(0 kJ)] = [(-1180.5 kJ) + (-1143.2 kJ)] - [(-103.85 kJ) + (0 kJ)] = (-2323.7 kJ) - (-103.85 kJ) = -2219.9 kJ

  11. Remember If a reaction is multiplied or divided by a number (whole or fractions), the ΔH must also be multiplied or divided by that same number

  12. Sample exercise 5.11 • Balanced equation: • H = nHf(products) - mHf(reactants) Calculate the standard enthalpy change for the combustion of 1 mol of benzene, C6H6(l), to CO2(g) and H2O(l). Compare the quantity of heat produced by 1.00 g of propane (-2220 kJ/mol) that that produced by 1.00 g of benzene.

  13. Practice exercise Using the standard enthalpies of formation listed in Table 5.3, calculate the enthalpy change for the combustion of 1 mol of ethanol: C2H5OH(l) + 3 O2(g) 2 CO2(g) + 3 H2O(l)

  14. Foods and Fuels Most of the fuel in the food we eat comes from carbohydrates and fats.

  15. Foods and Fuels The vast majority of the energy consumed in this country comes from fossil fuels.

  16. Problem solving – chapter 5

  17. Chapter 6 • If you look at the chapter alignments in the beginning of your book, section 6.5 is the section of this chapter we need to focus on. We’re going to go through a very brief overview of the chapter and then you’ll read the chapter, complete the overview, try some problems, and bring your questions back. • There’s a lot of crazy stuff in here, but don’t let it bog you down. The most important parts of this chapter are quantum numbers and electron configurations.

  18. Two main equations • Speed of light • c = λν • What is the value of c? • Energy of a photon • E = h ν • Remember that you can substitute one into the other!

  19. Electron configurations • What do you remember about electron configurations? • How do they relate to the periodic table?

  20. Rules for filling orbitals • Aufbau principle • Lowest energy orbital filled first • Pauli exclusion principle • An orbital can hold a maximum of 2 electrons and they must have opposite spin • Hund’s rule • Same spin first, then pair

  21. Electron Configurations • Distribution of all electrons in an atom • Consist of • Number denoting the energy level

  22. Electron Configurations • Distribution of all electrons in an atom • Consist of • Number denoting the energy level • Letter denoting the type of orbital

  23. Electron Configurations • Distribution of all electrons in an atom. • Consist of • Number denoting the energy level. • Letter denoting the type of orbital. • Superscript denoting the number of electrons in those orbitals.

  24. Orbital Diagrams • Each box represents one orbital. • Half-arrows represent the electrons. • The direction of the arrow represents the spin of the electron.

  25. Hund’s Rule “For degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”

  26. Periodic Table • We fill orbitals in increasing order of energy. • Different blocks on the periodic table, then correspond to different types of orbitals.

  27. Some Anomalies Some irregularities occur when there are enough electrons to half-fill s and d orbitals on a given row.

  28. Some Anomalies For instance, the electron configuration for copper is [Ar] 4s13d10 rather than the expected [Ar] 4s23d9.

  29. Some Anomalies • This occurs because the 4s and 3d orbitals are very close in energy. • These anomalies occur in f-block atoms, as well.

  30. Bell work • Turn in the chapter 6 overview. Scale of 1-10, how was chapter 6? Which topics were new? Which topics were review? • Draw the orbital diagram for the electron configuration of oxygen, atomic number 8. How many unpaired electrons does an oxygen atom possess?

  31. Agenda • Bell work • Chapter 6 – main points –quantum numbers • Homework (due Thursday, Oct 14) • Chapter 6 # 5, 8, 16, 18, 28, 36 (hint section 6.3 Rydberg’s equation), 50, 52, 54, 64, 68, 70, 72, 74, 100 • Chapter 4 and 5 test on Tuesday, Oct. 12

  32. Looking ahead… • Thursday (today) • chapter 6 problem solving • Homework – chapter 6 problems • Tuesday, September 27 • Test over chapters 4 and 5; lab on Enthalpy of Reaction • Homework: Chapter 7 overview, finish chap 6 problems • Thursday, Sept 29 • Introduce chapter 7, midterm talk • Due: Chapter 7 overview and problems from chapter 6

  33. Practice • Write the electron configuration for phosphorus, element 15. • How many unpaired electrons does a phosphorus atom possess?

  34. Sample 6.8 • What is the characteristic valence electron configuration of the group 7A elements, the halogens?

  35. Practice exercise • Which family of elements is characterized by an ns2np2 electron configuration in the outermost occupied shell?

  36. Sample 6.9 • Write the electron configuration for bismuth, element 83. • Write the condensed electron configuration for this element. • How many unpaired electrons does each atom of bismuth possess?

  37. Practice exercise • Use the periodic table to write the condensed electron configurations for (a) Co (atomic number 27), (b) Te (atomic number 52).

  38. Quantum Mechanics • Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. • It is known as quantum mechanics.

  39. Quantum Mechanics • The wave equation is designated with a lower case Greek psi (). • The square of the wave equation, 2, gives a probability density map of where an electron has a certain statistical likelihood of being at any given instant in time.

  40. Quantum Numbers • Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. • Each orbital describes a spatial distribution of electron density. • An orbital is described by a set of three quantum numbers.

  41. Quantum numbers • There are 4 quantum numbers • n • l • ml • ms

  42. Quantum numbers • They tell you which electron you are talking about/where the electron you are talking about is

  43. Quantum numbers • There are 4 quantum numbers • n – principle quantum number • l - Azimuthal quantum number • ml - magnetic quantum number • ms – spin quantum number

  44. Quantum numbers • There are 4 quantum numbers • n – energy level • l - shape • ml - # of subshells in sublevel • ms – ± ½ (spin up or down)

  45. Allowed values • n : any positive integer (1, 2, 3, …) • l : 0 to n-1 • ml : - l to l (including zero) • ms: ± ½ (spin up or down)

  46. Questions based on this… • Which set of quantum numbers is not allowed? (based on the allowed values) • Which set of quantum numbers is allowed? • The valence electrons of ___ can be represented by which set of quantum numbers:

  47. Principal Quantum Number, n • The principal quantum number, n, describes the energy level on which the orbital resides. • The values of n are integers ≥ 0. • Tells you the period

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