Chapters 7 & 8

1. Chapters 7 & 8 Quantum Theory and Periodic Relationships

2. Properties of Waves • Vibration • Transmits energy • Wavelength (λ) • Distance between • Frequency (υ) • Number/time, Hz or s-1 • Amplitude (A) • “Height” • Speed (u) • u= (λ) (υ) • How do we describe a wave?

3. Properties of Waves • Electrical and magnetic that vibrate perpendicular to each other • All the properties of a wave but can also be thought of as a stream of particles • Ex. Light particles are called photons • What are the components of an EM wave? • What are the properties of EM radiation?

4. Electromagnetic Waves • In a vacuum: • 3.0 x 108 m/s or c • The speed of light • Otherwise, it depends on the medium (substance) • We use the formula: • ν = c/λ • How fast do EM waves travel?

5. Wave Examples • Example 7.2 (pg. 246-7) • The wavelength of green light from a traffic signal is 522 nm. What is the frequency of this radiation? • ν = c/λ and c = 3.00 x 108 m/s • 5.75 x 1014 Hz • What is the wavelength (in meters) of an electromagnetic wave whose frequency is 3.64 x 107 Hz? • ν = c/λ and c = 3.00 x 108m/s • 8.24 m

6. Electromagnetic Sprectrum

7. Electromagnetic Spectrum • As wavelength increases, frequency decreases • Indirect relationship • About 5%, not much • It is called visible light • What is the relationship between wavelength and frequency? • Approximately what percent of the EM spectrum can we see?

8. Visible Light

9. Examples of EM Radiation • Radio • TV broadcasting • AM and FM broadcast radio • Avalanche beacons • Heart rate monitors • Cell phones • What are some common examples of non-visible EM energy?

10. Examples of EM Radiation • Microwave • Microwave ovens • Bluetooth headsets • Wireless internet • Radar • GPS • What are some common examples of non-visible EM energy?

11. Examples of EM Radiation • Infrared • Night vision goggles • Remote controls • Heat-seeking missiles • Heat lamp • What are some common examples of non-visible EM energy?

12. Examples of EM Radiation • Ultraviolet • Black lights • Sterilizing medical equipment • Water disinfection • Security images on money • Tanning beds • What are some common examples of non-visible EM energy?

13. Examples of EM Radiation • X-rays • Medical imaging • Airport security • Inspecting industrial welds • What are some common examples of non-visible EM energy?

14. Examples of EM Radiation • Gamma • Food irradiation • Cancer treatment • Soil density • Treating wood flooring • What are some common examples of non-visible EM energy?

15. Electromagnetic Spectrum • Percent of each type that reaches the surface.

16. Electromagnetic Spectrum

17. Planck’s Quantum Theory • Need Planck’s constant: • h = 6.63 x 10-34 Js • E = hυ • E represents a single quanta of energy • Energies of a wave are quantized • How can we calculate the energy of the wave?

18. Planck’s Quantum Theory • Energy can only be released or absorbed in multiples of hυ • Ex. 2hυ, 3hυ… • NOT 1.5hυ • These very specific energies are an atom’s electron orbitals • Convergence • What does quantized mean?

19. Energy Practice • Example 7.3 (pg. 249): Calculate the energy of a photon (light particle) with a wavelength of 5.00 x 104 nm. • h = 6.63 x 10-34Js • c = 3.00 x 108m/s • E = hυand c= λυ • Combined: E = hc/ λ • Answer = 3.98 x 10-21 J

20. Energy Practice • Example 7.3 (pg. 249): The energy of a photon is 5.87 x 10-20 J. What is its wavelength (in nanometers)? • h = 6.63 x 10-34Js • c = 3.00 x 108m/s • E = hυand c= λυ • Combined: E = hc/λ • Answer: 3.39 x 103 nm

21. More Energy Practice • Barcode scanners use a red light with a wavelength of 633 nm. Determine the energy of a photon from this light. • h = 6.63 x 10-34Js • c = 3.00 x 108 m/s • E = hυand c= λυ

22. More Energy Practice • Calculate the energy of a single photon of red light with a wavelength of 700.0 nm and then the energy of a mole of these photons. • h = 6.63 x 10-34Js • c = 3.00 x 108m/s • E = hυand c= λυ

23. More Energy Practice • What is the energy of cell phone radiation with a 1m wavelength? • h = 6.63 x 10-34Js • c = 3.00 x 108 m/s • E = hυand c= λυ

24. Brief History of the Atom • How did the current quantum model of an atom come to be?

25. Bohr’s Model of an Atom • Determine the element • Determine the total number of electrons • Place the electrons in the circular orbital’s following the 2.8.8.2 pattern • How do you arrange electrons in a Bohr model?

26. Bohr’s Model of an Atom • Neutral carbon has 6 electrons • Draw a circular diagram with the following: • 1st orbital – 2 • 2nd orbital - 4 (not full) • Principal quantum numbers (n) • 1st orbital → n=1 • What does carbon’s model look like? • What’s another way to name the different orbitals?

27. Bohr’s Model of an Atom • Atoms “like” to have 8 electrons in the n=2 and n=3 shells (orbitals) • n=1 is too small and can only hold 2 • Atoms will react (gain/lose electrons) until they have a full outer shell • The electrons in the outermost orbital • What is the octet rule? • What are valence electrons?

28. Bohr’s Model of an Atom • Please draw the appropriate Bohr model for: • Li • N • Ne • H • Mg2+ • S2- • Cl1- • Cl7+

29. Bohr’s Model of an Atom • Each orbital represents a specific quantized energy • As electrons move between them they absorb and release EM radiation specific to the energy difference between the orbitals • What is special about each orbital?

30. Emission Spectra • Identifies each element/compound with a unique light signature • Electrons first need to be excited by heat, light or electricity • Excited electrons will relax and release light at very specific λ’s based on the move • Ground vs. Excited State • What does an emission spectrum tell us?

31. E = hn E = hn

32. Visible Light and Emission Spectra

34. Quantum Numbers • Electron configuration - describing the distribution of electrons • Principal quantum number (n) • n = 1,2,3… • The orbital number • What is the purpose of quantum numbers? • What are the four quantum numbers?

35. Quantum Numbers • Angular momentum quantum number (l) • Tells us the shape • l = 0 to n -1 • What are the four quantum numbers?

37. Quantum Numbers • Magnetic quantum number (ml) • Tells us the orientation (how many subshells) • ml = -lto +l • Electron spin quantum number (ms) • Electron spin direction • Two spots (+ and –) • What are the four quantum numbers?

38. Orbital Energy Level Diagrams

39. Filling the Subshells • Hund’s Rule - Every orbital in a subshell is occupied with one electron before another is added • Pauli Exclusion Principle – No two electrons in an atom can have the same quantum numbers • What two rules must we follow when configuring electrons?

40. Learn by doing! • Determine the electron configuration for Ne.

41. Block We can see that one way the periodic table is arranged is by electron configuration! For example, Groups 1 and 2 are called s-block

42. Learn by doing! • Draw a Bohr diagram and determine the electron configuration for B. • Draw a Bohr diagram and determine the electron configuration for Mg. • Draw a Bohr diagram and determine the electron configuration for Ca.

43. A Brief History of the Periodic Table

44. A Brief History of the Periodic Table

45. Groups and Periods

46. IB Focus • The following groups and periods are the main focus in the IB curriculum:

47. Trends In Physical Properties • A general pattern across a period or down a group • Atomic number • Atomic radii • Ionic radii • 1st Ionization energy • Electronegativity • Melting points • What is a periodic trend? • What physical properties can we use to find trends?

48. Atomic Number • Number of protons within an atom • General trends: • Group (down) • Increases • Period (across) • Increases • What is the atomic number?

49. Atomic Radius • Half the distance from the center of an atom to the center of an adjacent atom • General trends: • Group (down) • Increases • Period (across) • Decreases • What is the atomic radius?

50. Atomic Radius • Group (down) – adding energy levels increases the radius • Period (across) – decreases across because effective nuclear chargeincreases (adding one more proton) • Why?