1. Chapter 7 Atomic Structure

3. Niels Bohr 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.

4. 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, but not because they are circling like planets.

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

6. Hydrogen spectrum n Emission spectrum because these are the colors it gives off or emits. n Called a line spectrum. n Developed the quantum model of the hydrogen atom. n There are just a few discrete lines showing 410 nm 434 nm 486 nm 656 nm

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

8. Kinds of EM waves n Light is made up of electromagnetic radiation n There are many types different (wavelength) and  different (wavelength) and  Frequency) n Light is only the part our eyes can detect. Gamma Rays Radio waves

10. Waves n n Waves have 3 primary characteristics: n n 1.Wavelength: distance between two peaks in a wave. n n 2.Frequency: number of waves per second that pass a given point in space. 3.Speed: speed of light is 3.0  10 8 m/s.

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

13. Wavelength and frequency can be interconverted. c  =  = frequency (s  1, Hz, cyc/s, or waves/s ) = wavelength (m) n n c = speed of light (m/s) in a vacuum is 3.00 x 10 8 m/s

14. The speed of light 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?

15. n 509.34M n 6.2 14

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

18. n Only certain energies are allowed for the hydrogen atom. n Energy in the atom is quantized.

19. Planck’s Constant  E = change in energy, in J h = Planck’s constant, 6.63  10  34 J s = frequency, in s  1 = wavelength, in m n n refrence table.pdf refrence table.pdf Transfer of energy is quantized, and can only occur in discrete units, called quanta.

20. When you want the change.  E = E final - E initial  E = E final - E initial

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

22. What is the wavelength? n of an electron with a mass of 9.11 x 10 -31 kg traveling at 5.31 x 10 6 m/s? n

23. n Electron Configuration n Na 2-8-1

25. 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

26. n Write the electron configuration for C, Cu Identify the following atom. 1s 2 2s 2 2p 5 3s 2 3p 3

27. Details n Hund’s Rule- The lowest energy configuration for an atom is the one have the maximum number of of unpaired electrons in the orbital. n C 1s 2 2s 2 2p 2

28. 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.

29. Exceptions n Cr = [Ar] 4s 1 3d 5 n Half filled orbitals. n Scientists aren’t sure of why it happens n same for Cu [Ar] 4s 1 3d 10 Draw the Orbital diagrams following Hund’s rule.

30. Ion Configurations n n To form cations from elements remove 1 or more e- from subshell of highest n n n P [Ne] 3s 2 3p 3 - 3e- ---> P 3+ [Ne] 3s 2 3p 0

31. Ion Configurations n n For transition metals, remove ns electrons and then (n - 1) electrons. n n Fe [Ar] 4s 2 3d 6 n n loses 2 electrons ---> Fe 2+ [Ar] 4s 0 3d 6 loses 3 electrons ---> Fe 3+ [Ar] 4s o 3d 5

32. Paramagnetic vs. Diamagnetic S Ca S Ca n Orbital notation for C Paramagnetic - Attracted by a magnet because of unpaired electrons Diamagnetic - Weakly repelled by magnet because of paired electrons

33. The Quantum Mechanical Model n A totally new approach. n De Broglie said matter could be like a wave.

34. n There are only certain allowed waves. n In the atom there are certain allowed waves called electrons. n 1925 Erwin Schroedinger described the wave function of the electron. n A lot of math but what is important are the solution.

35. Schroedinger’s Equation 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.

36. 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.

37. 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 of an orbital. n Has integer values >0

38. Quantum numbers Angular momentum quantum number l. Angular momentum quantum number l. shape of the orbital n integer values from 0 to n-1 l = 0 can hold 2 e- l = 0 can hold 2 e- l = 1 can hold 6 e- l = 1 can hold 6 e- l =2 can hold 10e- l =2 can hold 10e- l =3 can hold 14e- l =3 can hold 14e- l =4 can hold 18e- l =4 can hold 18e-

39. 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 Orientation of the orbital to each other Electron spin quantum number (m s ) Electron spin quantum number (m s ) n Can have 2 values. n either +1/2 or -1/2 n Spin of an electron

40. Since only 2 values of ms are allowed thefore an orbital can hold no more then two electrons n Pauli Exclusion Principle n In a given atom no two electrons can have the same set of 4 quantum #s.

41. n N=3 l =1 ml= -2 Ms =+1/2 n 2,3,-1,+1/2 n 2,1,0,-1/2 3,4,-1,+1/2 3,4,-1,+1/24,3,-3,+1/21,0,0,-1/2

42. L0= S orbitals

43. L1= P Orbitals

44. P orbitals

45. L2= D orbitals

46. L3= F orbitals

47. F orbitals

48. The Periodic Table n Developed independently by German Julius Lothar Meyer and Russian Dmitri Mendeleev (1870”s). n Put in columns by similar properties. n Predicted properties of missing elements.

49. Information Contained in the Periodic Table n n Each group member has the same valence electron configuration (these electrons primarily determine an atom’s chemistry). n n Certain groups have special names (alkali metals, halogens, ). n n Metals and nonmetals are characterized by their chemical and physical properties.

50. Broad Periodic Table Classifications n n Representative Elements (main group): filling s and p orbitals (Na, Al, Ne, O) n n Transition Elements: filling d orbitals (Fe, Co, Ni) n n Lanthanide and Actinide Series (inner transition elements): filling 4f and 5f orbitals (Eu, Am, Es) n

51. General Periodic Trends n n Atomic Radii n Ionization energy n n Electron Affinity n n Electronegativity

52. Reasons for General Periodic Trends Higher Z*. Larger orbitals. More sheilding.

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 n If shielding were completely effective, Z eff = 1

54. Atomic Radii

55. Ion Sizes n Does the size go up or down when gaining an electron to form an anion?

56. Ion Sizes n n ANIONS are LARGER than the atoms from which they come. n n The electron/proton attraction has gone DOWN and so size INCREASES. n n Trends in ion sizes are the same as atom sizes. Forming an anion.

57. Ion Sizes n Does the size go up or down when losing an electron to form a cation? Does the size go up or down when losing an electron to form a cation?

58. Ion Sizes n n. CATIONS are SMALLER than the atoms from which they come. n n The electron/proton attraction has gone UP and so size DECREASES. Forming a cation.

59. Periodic Trends n Ionization energy the energy required to remove an electron from a gaseous atom n Highest energy( the outer most) 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.

60. Ionization Energy n n IE = energy required to remove an electron from an atom in the gas phase. n n Mg (g) + 735 kJ ---> Mg + (g) + e-

61. Ionization Energy n n IE = energy required to remove an electron from an atom in the gas phase. n n Mg (g) + 735 kJ ---> Mg + (g) + e- n n Mg + (g) + 1445 kJ ---> Mg 2+ (g) + e-

62. 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.

63. 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, more shielding As you go down a group I 1 decreases because electrons are farther away, more shielding

64. n The first ionization energy for phosphorous is 1060 kj/mol and sulfur is 1005kj/mol. Why?

65. n Which atom has the largest first ionization energy,and which one has the smallest second ionization energy ? Explain your choice. n Ne n Na n Mg

66. 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

67. Electron Affinity n n A few elements GAIN electrons to form anions. n n Electron affinity is the energy involved with the addition of an electron in the gaseous state A(g) + e - ---> A - (g) E.A. =  E

68. Trends in Electron Affinity n n Affinity for electron increases across a period (EA becomes more negative). n n Affinity decreases down a group (EA becomes less negative).

69. Increasing Periodic Trends electronegativity, ionization energy, electron affinity atomic radii ionic & atomic radii ionization energy, electron affinity, & electronegativity

70. 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.

71. First Ionization energy Atomic number

72. n In the photoelectric effect, electrons are emitted from solids, liquids or gases when they absorb energy from light. Electrons emitted in this manner may be called photoelectrons. electrons

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