The Bohr Atom

1. The Bohr Atom • Thompson: plum pudding • Rutherford: nucleus in a sea of electrons. • Bohr: planetary model. • Wave velocity = (wavelength)(frequency). • Review of wave mechanics.

2. The Bohr Atom • The electromagnetic spectra: figure 6.2, p. 132 • The type of wave is determined by its wavelength and frequency. • Speed of light = 2.998 x 108 m/s

3. The Bohr Atom • A light wave has a frequency of 1.74 x 1017 hertz. What is its wavelength? What type of light is it? • Wavelength = 1.72 x 10-9m • (x ray)

4. Planck’s Hypothesis • Light is given off in bundles of energy called quanta. • E = h • Energy = planck’s constant * wave frequency • h = 6.626 x 10-34 J/hz

5. Energy Problems • A photon of light is found to have 3.63 x 10-22 Joules of energy. What is the wavelength of this light wave?  = 5.47 x 10-4 m

6. Energy Problems • A photon of light is found to have a wavelength of 5.00 x 102 nm. What is the energy of a photon of this light wave? E = 3.98 x 10-19 J

7. Planck’s Hypothesis • Atomic Spectra: produced when an electron moves from a higher to lower energy level, giving off light in the process.

8. Planck’s Hypothesis • Atomic Spectra: produced when an electron moves from a higher to lower energy level, giving off light in the process. E = Ehi - Elo = h = hc/

9. Planck’s Hypothesis • Ex. For the yellow line in the sodium spectra ( = 589.0 nm), find its frequency, quantum energy, and the energy released by one mol of sodium electrons.What is the energy difference between two energy levels of Na?

10. Planck’s Hypothesis 5.090 x 1014s-1 E = 3.373 x 10-19 J For one mol of electrons: E = 203.1 kJ Hence a two energy level difference = 203.1 kJ/mol

11. Bohr Model • Bohr postulated that an electron moves about the nucleus in a circular orbit of a fixed radius.

12. Bohr Model • The emission spectra of hydrogen:hydrogen absorbs energy when excited then gives it off when it returns to its ground state.

13. Bohr Model • The ground state of an electron represents its lowest orbit. The excited state represents any other possible orbit.

14. Bohr Model • Hydrogen only emits this absorbed energy at certain visible wavelengths. Bohr reasoned this related to certain allowed electron orbits.

15. Bohr Model • To calculate the energy of an allowed energy level: En = (-2.180 x 10-18 J)/n2, where n = 1, 2, 3, … • RH = Rydberg constant = 2.180 x 10-18 J

16. Bohr Model • In the Bohr atom, calculate the energy released as an electron moves from the third to the second energy level. What is the wavelength of the emitted light?

17. Bohr Model • E3 = -2.422 x 10-19 J; E2 = -5.450 x 10-19 J • Ehi - Elo = 3.028 x 10-19 J = hc/E = 6.560 x 10-7 m = 656.0 nm (how does this compare to the Balmer series?)

18. Modern Atomic Structure • Bohr: emission spectra turned out to be several lines at each level, not singular lines. • De Broglie: wavelengths can be predicted based on the mass and velocity of a particle.

19. Modern Atomic Structure • Wave/particle duality. • Planck - waves can act like particles E=hn • DeBroglie - hey… then particles can act as waves, mc2 = E = hn, l=h/mv.

20. Modern Atomic Structure • Wave/particle duality. • Experiments can only demonstrate one of these qualities at a time. • Particle behavior: photoelectric effect (solar powered calculator).

21. Modern Atomic Structure • Wave/particle duality. • Wave behavior: refraction (changes speed in different media), defraction (bends around barriers), reflection

22. Modern Atomic Structure • Heisenberg Uncertainty Principle: both the momentum and position of a particle can not be precisely known at the same time.

23. Modern Atomic Structure • Therefore, we can only refer to the probability of finding an electron in a region; we cannot specify the path.

24. Modern Atomic Structure • Schrodinger: wave equations (y2) can be used to predict the region of probability for locating an electron.

25. Modern Atomic Structure • An Electron moves at high velocities usually on the surface of this region. • An electron effectively fills the surface. (fan analogy)

26. Modern Atomic Structure • Quantum numbers are used to describe the location of electrons in atoms. • Importance: model of atoms and bonding theory.

27. Modern Atomic Structure • Principle Quantum Number, n: energy level. • The higher the number the larger the region. • Corresponds to the periodic table.

28. Modern Atomic Structure • Principle energy level: the value of n is the main factor that determines the energy of an electron and its distance from the nucleus. • Maximum electron capacity of a level = 2n2

29. Modern Atomic Structure • Second Quantum Number, l: refers to energy sublevels. • The number of sublevels equals the principal quantum number.

30. Modern Atomic Structure • Sublevels do not have the same energy. • Sublevels from one principal level can overlap sublevels from another. Figure 6.7, p. 151

31. Modern Atomic Structure • 3rd Quantum Number, m: refers to the orientation of the suborbital. • s,p,d and f orbitals. • Degenerate orbitals (geometries and orientations - capacities)

32. Modern Atomic Structure • Each orbital has the capacity of two electrons. The s orbitals are spherically symmetric about the nucleus; p orbitals are dumbell shaped and at right angles to each other.

33. Modern Atomic Structure • 4th Quantum Number: refers to the spin. • ms = 1/2 or -1/2

34. Modern Atomic Structure • Pauli Exclusion Principle: each electron can be described by a unique set of 4 quantum numbers.

35. Modern Atomic Structure n = primary energy level l = sublevels • 0 = s-orbital • 1 = p-orbital • 2 = d-orbital.

36. Modern Atomic Structure • m = orientation of the orbital ( -l to +l) • ex. p-orbital • px = -1 • py = 0 • pz = +1

37. Modern Atomic Structure • spin = +1/2 or -1/2 • ex: • 1st electron = 1 0 0 +1/2. • 2nd electron = 1 0 0 -1/2 • 3rd electron = 2 0 0 +1/2

38. Modern Atomic Structure • Hund’s Rule: electrons fill unoccupied degenerate orbitals before pairing. • Find the quantum numbers for the 5th and 7th electrons.

39. Modern Atomic Structure • Basic electron configurations. • Orbital diagrams. • Core configurations.

40. Modern Atomic Structure • Note: 1. 2 e- in an orbital have opposed spins. • 2. When several orbitals of the same sublevel are available, electrons enter one at a time with parallel spins.

41. Modern Atomic Structure • Pneumonic device for remembering the filling order.

42. Modern Atomic Structure • Use the periodic table to locate the outermost electrons of an atom. • S-block and p-block elements match the row.

43. The Transition Metals • d-block elements (groups 3-12). • Energy sublevel overlap: ex - 4s vs. 3d • Multiple valence

44. The Periodic Table • Brightly colored compounds and solutions.

45. The Periodic Table • Lanthanoids: elements 57-70, begins 4f block. • Actinoids: elements 89-102, begins 5f block. • Rare earth metals.

46. The Periodic Table • Nature tends towards stability. • Atoms seek bonding situations that result in stable electron configurations (ex. Share electrons).

47. The Periodic Table • Octet rule: eight electrons in the outer level (s & p’s?) render an atom unreactive. • Atoms seek to lose, gain, or share electrons to seek a stable octet of electrons.

48. The Periodic Table • An atom having a filled or half-filled sublevel is slightly more stable than an atom without. • Full sublevels are more stable than half-filled.

49. The Periodic Table • Full outer levels are more stable than full sublevels.

50. The Periodic Table • Electron promotion: an electron can be promoted to a slightly higher sublevel in order to produce a full or half filled sublevel. • Ex - Cr and Cu (p. 154)