Chapter 5 – Models of the Atom

2. Chapter 5 – Models of the Atom

3. Friday, October 1 • Do Now – Current Events and Chemistry • Objective – Chapter 4 Review 5.1 Models of the Atom • Homework – Start Pg. 132 # 1-6 (Due Tuesday)

4. Supersonic Man How does this relate to anything we have talked about already this year? Write you idea in your notes

5. 10/11/11 • Objective: to explore the electronic structure of the atom • Do now: Take out HW questions (Pg 119 # 25-28, 30-33)

6. Review Questions – Pg 119 # 25-28, 30-33 25) Atoms of different elements contain different numbers of protons 26) Mass # - Atomic # = # of neutrons 27) Isotopes have different mass numbers and different numbers of neutrons 28) For each isotope, multiply its atomic mass by its % abundance, and then add the products • 194 is the mass number • The atomic mass is a weighted average of the masses of its isotopes • Lithium-6, 3p+, 3 e-, 3 no; Lithium-7, 3p+, 3 e-, 4 no; Calcium-42, 20 p+, 20 e-, 22 no; Calcium-44, 20 p+, 20 e-, 24 no; Selenium-80, 34 p+, 34 e-, 46 no; Selenium-78, 34 p+, 34 e-, 44 no; • Beryllium, Magnesium, Strontium, Barium, Radium

8. Review what we know Democritus Dalton Thompson Rutherford

9. The Bohr Model • Rutherford model: electrons move around the nucleus • Bohr: How do they move around the nucleus?

10. Bohr model • Bohr knew that different elements could emit different colors of light when heated • Each element has its own specific color • Why does light come in different colors? Iron Lithium

11. 1) Rutherford’s atomic model needed to be replaced because it could not explain why elements give off characteristic colors • 2) In the bohr model, an electron is found only in specific circular paths around the nucleus. • 3) the QMM determines the allowed energy levels an electron can have and the likelihood of finding an electron in a given location • 4) The sublevels have different shapes • 5) An electron can move from one energy level to another by gaining or losing energy • 6) To move from one energy level to another requires emission or absorption of an exact amount of energy.

12. Light! • Light is composed of waves • Shorter wavelength = higher energy • Blue = short wavelength • Long wavelength = lower energy • Red = long wavelength

13. Bohr model • Bohr put this knowledge together to explain how electrons are organized in atoms • Electrons exist in fixed energy levels • They can jump from one level to another • But they are never in between two levels

14. Bohr Model • Energy levels • All electrons in a given energy level have equal energy • Electrons in higher energy levels have greater energy

15. Bohr Model • Energy has to be added for an electron to level up • Energy has to be released for an electron to level down • This energy is released as light

16. Bohr Model • So why do different elements produce different colors of light? • Rydberg equation: λ = wavelength R = 0.0110 nm−1 (Rydberg constant) m = ending energy level (1, 2, 3…) n = starting energy level ( n>m) • Each element’s electrons are excited to higher energy levels in different patterns that produce different wavelengths of light

17. Emission Spectra • Each element emits a combination of wavelengths that we interpret as a unique color

18. 10/18/11 • Do now: How are electrons localized around the nucleus in the Bohr model? • What sort of regions do they move in? • How do they move? • How many electrons are in each region? • How does an electron’s position relate to its energy? • Do later: Complete your orbital poster

19. Bohr Model v. Quantum Mechanical Model • In the Bohr model, electrons travelled in circular orbits around the nucleus • 2 in first shell, 8 in second shell, 8 in third shell… • One shell for each row of the periodic table • All electrons in a shell have equal energy

20. Quantum Mechanical Model • Our understanding of energy levels has improved since the Bohr model • Electrons do not orbit around the nucleus in circles. • They exist in 3-D orbitals – regions of space where you are likely to find an electron • Orbitals show the probability of where an electron is likely to be. • (this is where chemistry starts to get weird)

21. Where is the propeller?

22. The Quantum Mechanical Model • The QMM is based on the Schrodinger equation • Predicts where electrons are likely to be in an atom based on their energy • The solutions to this equation are orbitals

23. Atomic Orbitals • Atomic Orbitals are regions of space around a nucleus where a given electron is likely (90%) to be. • 4 kinds of orbitals: s, p, d, f • Each has a different shape • Each orbital can hold up to 2 electrons

25. Bohr Model v. Quantum Mechanical Model • In the Bohr model, electrons existed in shells around the nucleus • 2 in first shell, 8 in second shell, 8 in third shell… • One shell for each row of the periodic table

26. Bohr Model v. Quantum Mechanical Model • In the QMM, electrons are found in energy levels that are subdivided into orbitals • Each orbital can hold up to 2 electrons • 1st energy level has 1 orbital (2 electrons) • 2nd energy level has 4 orbitals (8 electrons) • 3rd energy level has 9 orbitals (18 electrons)

27. S Orbitals • The first energy level only has 1 orbital • 1s orbital • Sphere of space where up to 2 electrons are likely to be found

28. P orbitals • The second energy level has 1 2s orbital and 3 2p orbitals • P orbitals are shaped like barbells • Each holds up to 2 electrons, 6 e-s total • 3 p orbitals oriented on the x, y, and z axes

29. D orbitals • The third energy level has 1 3s orbital, 3 3p orbitals and 5 3d orbitals • 9 orbitals, 18 electrons total

30. F orbitals • The fourth energy level has 1 4s orbital, 3 4p orbitals, 5 4d orbitals, and 7 4f orbitals • 16 orbitals, 32 electrons total

31. Review of 5.1 Principal Number of Energy Levels Sublevels Orbitals 1 1 1s (1 orbital) 2 2 2s (1 orbital), 2p (3 orbitals) 3 3 3s (1 orbital,) 3p (3 orbitals), 3d (5 orbitals) 4 4 4s (1 orbital,) 4p (3 orbitals), 4d (5 orbitals), 4f (7 orbitals) Chlorine has 17 electrons 1s2 2s2 2p63s23p5

33. Energy of Orbitals • Two trends • More complex  higher energy • (f > d > p > s) • Higher energy level  higher energy • (p3 > p2 > p1)

34. Bohr Model v. Quantum Mechanical Model • Like the Bohr Model, in the QMM: • Electrons always fill in a lower energy orbital before a higher energy orbital • Higher energy orbitals are farther from the nucleus

35. Practice time! • Which orbital has higher energy electrons? • 2S or 1S • 2P or 2S • 3D or 2P • 3P or 4S

36. 10/20 • Objective: To practice our knowledge of the QMM and to understand why orbitals take their specific shapes • Do now: • Take out your orbital/energy level poster • Take 5 minutes to fill in the worksheet at the front desk • Do later: Read section 5.3, questions 16-21

38. Quantum Mechanical Model Review • Electrons are localized in 3-D regions called orbitals • 2 e- in each orbital • 4 kinds of orbitals: s, p, d, f • 1st energy level has 1 s1 orbital • 2nd energy level has 1 s2 & 3 p2 orbitals • 3rd energy level has 1 s3, 3 p3, & 5 d3 orbitals • 4th energy level has 1 s4, 3 p4, 5 d4, & 7 f4 orbitals

39. Quantum Mechanical Model Review • All electrons in a given orbital have equal energy

40. Quantum Mechanical Model Review • Within an energy level, more complex orbitals have higher energy electrons

41. Quantum Mechanical Model Review • Electrons in the same kind of orbital will have greater energy at higher energy levels

42. Practice time! • Which orbital has higher energy electrons? • 2S or 1S • 2P or 2S • 3D or 2P • 3P or 4S

43. Quantum Mechanical Model Review • Higher energy orbitals are located farther out from the nucleus

44. Practice problem! • If you were drawing 2p, 3p, and 4p orbitals, how would your drawings differ?

45. Orbital Shapes • Why do orbitals take these funny shapes? • Honors topic • Physics detour!

46. 10/21/11 • Objective: • To practice calculating wave measurements • To explore the wave-particle duality and its history • Do now: Check your homework with a partner • Do later: Interference worksheet

47. Homework • 16) Frequency and wavelength of light are inversely proportional to each other. λ=c/ν, ν=c/λ • 17) Electrons in atoms absorb energy as they move to energy levels, and then lose the energy by emitting it as light as they drop back • 18) The light emitted in an electronic transition from higher to lower energy levels has a frequency that is directly proportional to the energy change of the electron • 19)QM describes the motions of atoms and subatomic particles; classical mechanics describes the motion of larger bodies • 20)Electron transitions from higher levels to n=1 • 21) c, a, b

48. Waves • Waves are predictable disturbances that travel through space, usually accompanied by the transfer of energy • Transverse and longitudinal Amplitude Wavelength (λ) Frequency (v)

49. 10/24/11 • Objective: • To review the properties of waves • To explore the experiments that led to quantum mechanics • Do Now: Take 7 minutes to finish the waves practice sheet. Use a calculator. • Do Later: Worksheet – Experimental foundations of Quantum Mechanics