1. 1 Chapter 6 Atomic Structure

2. 2 Light n Made up of electromagnetic radiation. n Can behave as a particle as well as a wave

3. 3 Parts of a wave Wavelength Frequency = number of cycles in one second Measured in hertz 1 hertz = 1 cycle/second

4. 4 Frequency = Frequency =

5. 5 Kinds of EMR waves n There are radio, micro, IR, Vis, UV, X- rays and gamma rays Each have different and Each have different and n Visible light is only the part our eyes can detect. Gamma Rays Radio waves

6. 6 The speed of light n in a vacuum is 2.998 x 10 8 m/s n Denoted by c c = c = n What is the wavelength of light with a frequency 5.89 x 10 5 Hz? n What is the frequency of blue light with a wavelength of 484 nm?

7. 7 In 1900 n Matter and energy were seen as different from each other in fundamental ways. n Matter was particles. n Energy could come in waves, with any frequency. n Max Planck found that as the cooling of hot objects couldn’t be explained by viewing energy as a wave.

8. 8 Energy is Quantized Planck found  E came in chunks with size h Planck found  E came in chunks with size h  E = nh  E = nh n where n is an integer. n and h is Planck’s constant n h = 6.626 x 10 -34 J s these packets of h are called quantum these packets of h are called quantum

9. 9 Einstein is next n Said electromagnetic radiation is quantized in particles called photons. Each photon has energy (E) = h = hc/ Each photon has energy (E) = h = hc/ n Combine this with E = mc 2 n You get the apparent mass of a photon. m = h / ( c) m = h / ( c)

10. 10 Which is it? n Is energy a wave like light, or a particle? n Yes n Concept is called the Wave -Particle duality. n What about the other way, is matter a wave? n Yes

11. 11 Matter as a wave Using the velocity v instead of the wavelength we get. Using the velocity v instead of the wavelength we get. De Broglie’s equation = h/mv De Broglie’s equation = h/mv n Can calculate the wavelength of an object.

12. 12 Examples n The laser light of a CD is 7.80 x 10 2 m. What is the frequency of this light? n What is the energy of a photon of this light? n What is the apparent mass of a photon of this light? n What is the energy of a mole of these photons?

13. 13 Diffraction n When light passes through, or reflects off, a series of thinly spaced lines, it creates a rainbow effect n because the waves interfere with each other.

14. 14 A wave moves toward a slit.

15. 15 Comes out as a curve

16. 16 What will an electron do? n It has mass, so it is matter. n A particle can only go through one hole. n A wave through many holes. n An electron makes an interference pattern. n It behaves like a wave. n Other matter has wavelengths too small to notice.

17. 17 Spectrum n The range of frequencies present in light. n White light has a continuous spectrum. n All the colors are possible. n A rainbow.

18. 18 Hydrogen spectrum n Emission spectrum because these are the colors it gives off or emits. n Called a line spectrum. n There are just a few discrete lines showing 410 nm 434 nm 486 nm 656 nm

19. 19 What this means n Only certain energies are absorbed for the hydrogen atom. n So, it can only give off certain energies. Use  E = h  = hc / Use  E = h  = hc / n Energy in the atom is quantized.

20. 20 Niels Bohr n Developed the quantum model of the hydrogen atom. n He said the atom was like a solar system. n The electrons were attracted to the nucleus because of opposite charges. n Didn’t fall in to the nucleus because it was moving around.

21. 21 The Bohr Ring Atom n He didn’t know why but only certain energies were absorbed. n He called these allowed energies energy levels. n Putting energy into the atom moved the electron away from the nucleus. n From its ground state to its excited state. n When it returns to ground state it gives off light of a certain energy.

22. 22 The Bohr Ring Atom n = 3 n = 4 n = 2 n = 1

23. 23 The Bohr Model Equation n R H Rydberg constant -2.178 x 10 -18 J n n is the energy level n E = -2.178 x 10 -18 J / n 2 n n = 1 is called the ground state n Used to find the energy of one hydrogen electron

24. 24 We are worried about the change n When the electron moves from one energy level to another.  E = E final - E initial  E = E final - E initial n hv = -2.178 x 10 -18 J x (1/ n f 2 - 1/ n i 2 )

25. 25 Examples n Calculate the energy need to move an electron from its first to third energy levels. n Calculate the frequency released when an electron moves from n= 4 to n=2 in a hydrogen atom. n Calculate the energy released when an electron moves from n= 5 to n=3

26. 26 When is it true? n Only for hydrogen atoms and other monoelectronic species (He +1 ). n Why the negative sign? n B/c to increase the energy of the electron you make it closer to the nucleus.

27. 27 The Bohr Model n Doesn’t work. n Only works for hydrogen atoms. n Electrons don’t move in circles. n The quantization of energy is right

28. 28 The Quantum Mechanical Model n A totally new approach. n De Broglie said matter could be like a wave. n De Broglie said they were like standing waves. n The vibrations of a stringed instrument.

29. 29 Schroedinger’s Equation n The wave is a function of (x, y, z) n Solutions to the equation are called orbitals. n These are not Bohr orbits. n Each solution is tied to a certain energy. n These are the energy levels.

30. 30 There is a limit to what we can know n We can’t know how the electron is moving or how it gets from one energy level to another. n The Heisenberg Uncertainty Principle. n There is a limit to how well we can know both the position and the momentum of an object.

31. 31 What does the wave Function mean? n nothing. n it is not possible to visually map it. n The square of the function is the probability of finding an electron near a particular spot. n best way to visualize it is by mapping the places where the electron is likely to be found.

32. 32 Probability Distance from nucleus

33. 33 Defining the size n The nodal surface. n The size that encloses 90% to the total electron probability. n NOT at a certain distance, but a most likely distance. n For the first solution it is a sphere.

34. 34 Quantum Numbers n There are many solutions to Schroedinger’s equation n Each solution can be described with quantum numbers that describe some aspect of the solution. n Principal quantum number (n) size and energy of an orbital. n Has integer values >0

35. 35 Quantum numbers Angular momentum quantum number l. Angular momentum quantum number l. n shape of the orbital. n integer values from 0 to n-1 l = 0 is called s l = 0 is called s l = 1 is called p l = 1 is called p l =2 is called d l =2 is called d l =3 is called f l =3 is called f l =4 is called g l =4 is called g

36. 36 S orbitals

37. 37 P orbitals

38. 38 P Orbitals

39. 39 D orbitals

40. 40 F orbitals

41. 41 F orbitals

42. 42 Quantum numbers Magnetic quantum number (m l ) Magnetic quantum number (m l ) integer values between - l and + l integer values between - l and + l n tells direction of each shape Electron spin quantum number (m s ) Electron spin quantum number (m s ) n Can have 2 values. n either +1/2 or -1/2

43. 43 Polyelectronic Atoms n More than one electron. n three energy contributions. n The kinetic energy of moving electrons. n The potential energy of the attraction between the nucleus and the electrons. n The potential energy from repulsion of electrons.

44. 44 The Periodic Table n Developed independently by German Meyer and Russian Mendeleev (1870’s). n Didn’t know much about atom. n Put in columns by similar properties. n Predicted properties of missing elements.

45. 45 Aufbau Principle n Aufbau is German for building up. n As the protons are added one by one, the electrons fill up hydrogen- like orbitals. n Fill up in order of energy levels.

46. 46 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f He with 2 electrons

47. 47 Details n Valence electrons- the electrons in the outermost energy levels (not d or f). n Core electrons- the inner electrons. n Hund’s Rule- The lowest energy configuration for an atom is the one that has the maximum number of unpaired electrons in the orbital. n C = 1s 2 2s 2 2p 2

48. 48 Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2 electrons 2s 2 4 2p 6 3s 2 12 3p 6 4s 2 20 3d 10 4p 6 5s 2 38 4d 10 5p 6 6s 2 56

49. 49 Details n Elements in the same column have the same electron configuration. n Put in columns because of similar properties. n Similar properties because of electron configuration. n Noble gases have filled energy levels. n Transition metals are filling the d orbitals

50. 50 Exceptions n Cr = [Ar] 4s 1 3d 5 Mo as well n Half filled orbitals. n Scientists aren’t sure of why it happens n same for Cu [Ar] 4s 1 3d 10 and Ag

51. 51 More exceptions n Lanthanum La: [Xe] 6s 2 5d 1 n Cerium Ce: [Xe] 6s 2 4f 1 5d 1 n Promethium Pr: [Xe] 6s 2 4f 3 5d 0 n Gadolinium Gd: [Xe] 6s 2 4f 7 5d 1 n Lutetium Pr: [Xe] 6s 2 4f 14 5d 1 n We’ll just pretend that all except Cu and Cr follow the rules.

52. 52 n Ionization Energy- The energy necessary to remove an electron from a gaseous atom. n The ionization energy for a 1s electron from sodium is 1.39 x 10 5 kJ/mol. n The ionization energy for a 3s electron from sodium is 4.95 x 10 2 kJ/mol. n Demonstrates shielding.

53. 53 Shielding n Electrons on the higher energy levels tend to be farther out. n Have to look through the other electrons to see the nucleus. n They are less effected by the nucleus. n lower effective nuclear charge

54. 54 Penetration effect n The outer energy levels penetrate the inner levels so the shielding of the core electrons is not totally effective. n from most penetration to least penetration the order is n ns > np > nd > nf (within the same energy level). n This is what gives us our order of filling, electrons prefer s and p.

55. 55 How orbitals differ n The more positive the nucleus, the smaller the orbital. n A sodium 1s orbital is the same shape as a hydrogen 1s orbital, but it is smaller because the electron is more strongly attracted to the nucleus. n The helium 1s is smaller as well. n This provides for better shielding.

56. 56 Periodic Trends n Ionization energy the energy required to remove an electron form a gaseous atom n Highest energy electron removed first. First ionization energy ( I 1 ) is that required to remove the first electron. First ionization energy ( I 1 ) is that required to remove the first electron. Second ionization energy ( I 2 ) - the second electron Second ionization energy ( I 2 ) - the second electron n etc. etc.

57. 57 Trends in ionization energy n for Mg I 1 = 735 kJ/moleI 1 = 735 kJ/mole I 2 = 1445 kJ/moleI 2 = 1445 kJ/mole I 3 = 7730 kJ/moleI 3 = 7730 kJ/mole n The effective nuclear charge increases as you remove electrons. n It takes much more energy to remove a core electron than a valence electron because there is less shielding.

58. 58 Explain this trend n For Al I 1 = 580 kJ/moleI 1 = 580 kJ/mole I 2 = 1815 kJ/moleI 2 = 1815 kJ/mole I 3 = 2740 kJ/moleI 3 = 2740 kJ/mole I 4 = 11,600 kJ/moleI 4 = 11,600 kJ/mole

59. 59 Across a Period Generally from left to right, I 1 increases because Generally from left to right, I 1 increases because n there is a greater nuclear charge with the same shielding. As you go down a group I 1 decreases because electrons are farther away. As you go down a group I 1 decreases because electrons are farther away.

60. 60 It is not that simple n Half filled and filled orbitals are harder to remove electrons from. n here’s what it looks like on a graph.

61. 61 First Ionization energy Atomic number

62. 62 First Ionization energy Atomic number

63. 63 First Ionization energy Atomic number

64. 64 Atomic and Ionic radius n Atomic radius is the relative size of an atom –Decreases across a period –Increases down a group n Ionic radius is the relative size of an ion –Positive ions get smaller –Negative ions get larger –Effective charge

65. 65 Electronegativity n Is the tendency of atoms to gain electrons n Increases across n Decreases down n Scale of 4.0 (F) – 0.8 (Cs)

66. 66 Special groups n Group 1 Alkali metals n Group 2 Alkaline earth metals n Group 17 Halogens n Group 18 Noble gases