1. 1 Electrons in Atoms

2. 2 Have you even wondered why different atoms absorb and emit light of different colors? The transition of electrons within sublevels releases an amount of energy. If this energy corresponds to the visible section of the spectrum, then we observe colors.

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5. 5 Gamma rays is the most energetic radiation and shortest wavelength. Gamma rays are produced by the sun, by the stars, by some unstable atomic nuclei on earth. Human exposure to gamma rays are dangerous because of the high energy, they can damage biological molecules.

6. 6 X-rays can also damage biological molecules, but it requires excessive exposure to that radiation. Ultraviolet light is an important component of sunlight. It is not as energetic as gamma or X-rays, but still carries enough energy to damage biological molecules. Excessive exposure to ultraviolet light increases the risk of skin cancer and cataracts.

7. 7 Infrared light – the heat that you feel when you place your hand near a hot object is infrared light. All hot objects including the human body emit infrared light. It is invisible to our eyes, but infrared sensors can detect the infrared light and these sensors are often used in night vision to “see” in the dark. Visible light the only light detected by the human eye.

8. 8 Microwave light – lower energy but it is absorbed by water and therefore heat substances that contain water. For this reason substances that contain water, such as food, are warmed in a microwave oven, but substances that do not contain water such as a plate are not.

9. 9 The longest wavelengths, are used to transmit signals responsible for AM and FM radios, cellular phones, television and other forms of communication. Radiowaves

10. 10 Parts of a wave Wavelength Wavelength = distance between two consecutive crests. Measure in meters, cm, nm Recall 1 m = 1x10 9 nm Frequency = number of cycles in one second Measured in hertz 1 hertz = 1 /second

11. 11 Frequency = Frequency =

12. 12 The longer the wavelength (λ), the shorter the frequency (ν). The relationship is expressed as: c = ν λ where c = speed of light (constant) = 3.0x10 8 m/s

13. 13 What is the wavelength of light with a frequency 5.89 x 10 5 Hz?

14. 14 What is the frequency of blue light with a wavelength of 484 nm?

15. 15 Understanding the Spectrum

16. 16 The Hydrogen Line Emission Spectrum The lowest energy state of an atom is its ground state. A state in which an atom has the highest potential energy than it has in its ground state is an excited state.

17. 17 The Hydrogen Line Emission Spectrum Absorption and Emission Spectra Absorption and Emission Spectra http://www.flinnsci.com/atomicspectrum

18. 18 Energy is Quantized Planck found  E came in “chunks” with size h Planck found  E came in “chunks” with size h  E = h  E = h 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 n See Planck’s ideas See Planck’s ideas See Planck’s ideas

19. 19 A photon is a particle of electromagnetic radiation with zero mass and carrying a quantum of energy. The energy of a particular photon depends on the frequency of the radiation: E photon = h ν

20. 20 Niels Bohr n Developed the planetarium model of the atom. n He said the atom was like a solar system where electrons rotate around the nucleus like planets around the sun. n According to the model, the electron can circle the nucleus only in allowed paths or orbits. n The energy of the electron is higher when the electron is in orbits that are successively farther from the nucleus

21. 21 Bohr’s model

22. 22 The Bohr Model of the Atom  Energy must be added to an atom in order to move an electron from a lower energy level to a higher energy level. This is absorption.  When an electron fall to a lower energy level, a photon is emitted and the process is called emission.

23. 23 The Bohr Model of the Atom

24. 24 Visual concept

25. 25 Energy When energy is put into an atom, the absorbed energy allows the electron to reach higher energy levels. The electron will be in an excited state. n=1 n=5 n=4 n=3 n=2 n=6 n=7

26. 26 Energy Only certain energies are allowed, so in order for the electron to “jump” to a higher energy level, it will have to absorb the energy equal to the energy gap between the energy levels. n=1 n=5 n=4 n=3 n=2 n=6 n=7 ΔE=E 4 -E 1

27. 27 Energy When an excited electron drops to a lower energy level, the atom emits energy. The emitted energy is what makes the emission spectra. n=1 n=5 n=4 n=3 n=2 n=6 n=7

28. 28 Energy The energy released by the electron is directly proportional to the frequency (ν) and inversely proportional to the wavelength (λ)of the radiation. n=1 n=5 n=4 n=3 n=2 n=6 n=7

29. 29 Energy This transition will show a line in the spectra in a different position (wavelength) as the transition n=7  n=1 n=1 n=5 n=4 n=3 n=2 n=6 n=7

30. 30 The Quantum Mechanical Model n A totally new approach. Schröedinger’s Equation n Developed an equation that treated electrons in atoms as waves. n The wave function is a F(x, y, z) n Electrons do not travel around the nucleus in neat orbits, like Bohr postulated. n Instead, the exist in certain regions called orbitals. n Orbital is a three dimensional region around the nucleus that indicated the probable location of an electron.

31. 31 The Heisenberg Uncertainty Principle  It is impossible to determine simultaneously both the position and the velocity of an electron or any other particle.  Both, the Heisenberg Uncertainty Principle and the Schrödinger wave equation laid the foundation for modern quantum theory. Heisenberg Principle

32. 32 Quantum theory Quantum numbers and orbitals The solutions to the Schröndinger wave equation are called quantum numbers and they describe the probability of finding the electron around the nucleus.The solutions to the Schröndinger wave equation are called quantum numbers and they describe the probability of finding the electron around the nucleus.

33. 33 S orbital

34. 34 P orbitals

35. 35 D orbitals

36. 36 F orbitals

37. 37 F orbitals

38. 38 Electron spin quantum number (s)Electron spin quantum number (s) Can have 2 values.Can have 2 values. either +1/2 or -1/2either +1/2 or -1/2

39. 39 Energy level 1 Energy level 2 Energy level 3 Energy level 44 sublevels 3 sublevels 2 sublevels 1 sublevel 4f (14 electrons) 3p (6 electrons) 3s (2 electrons) 2p (6 electrons) 2s (2 electrons) 1s (2 electrons) 4s (2 electrons) 4p (6 electrons) 4d (10 electrons) 3d (10 electrons) 2(n) 2 2(1) 2 = 2 2(2) 2 = 8 2(3) 2 = 18 2(4) 2 = 32

40. 40 Aufbau Principle Aufbau Principle n Aufbau is German for building up. Electrons enter orbitals from low energy to high energy. The order of orbitals based on their energies is: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p …

41. 41 Pauli Exclusion Principle Pauli Exclusion Principle n No two electrons in the same atom can have the same set of quantum numbers. n Even if two electrons are located in the same energy level and same orbital, they must be different in the spin number.

42. 42 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f H with 1 electron

43. 43 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

44. 44 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Li with 3 electrons

45. 45 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Be with 4 electrons

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

47. 47 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f C with 6 electrons

48. 48 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f N with 7 electrons

49. 49 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f O with 8 electrons

50. 50 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f F with 9 electrons

51. 51 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Ne with 10 electrons

52. 52 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Na with 11 electrons

53. 53 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Mg with 12 electrons

54. 54 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Ar with 18 electrons

55. 55 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f K with 19 electrons

56. 56 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Ca with 20 electrons

57. 57 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Sc with 21 electrons

58. 58 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Mn with 25 electrons

59. 59 Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Ga with 31 electrons

60. 60 Orbital NotationOrbital Notation see PPT see PPT Orbital Notationsee PPT n Valence electrons- the electrons in the outermost energy levels (not d). n Core electrons- the inner electrons. n Hund’s Rule- The lowest energy configuration for an atom is the one have the maximum number of unpaired electrons in the orbital. n C 1s 2 2s 2 2p 2

61. 61 Orbital notation or box notation allows to consider Hund’s rule and the spin of the electron. For example : 8 O 1s 2 2s 2 2p 4 With orbital notation, the orbitals are represented by boxes (or some authors use circles). That way the s orbital will only be represented by 1 box, p orbitals by 3 boxes, and so on. 8O8O 1s 2 2s 2 2p 4 Oxygen is paramagnetic with 2 unpaired electrons

62. 62 25 Mn 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 Manganese is paramagnetic with five unpaired electrons. Paramagnetism is the property of substances of being attracted by an external magnetic field. It is present in all substances with at least 1 unpaired e -.

63. 63 30 Zn 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 Zinc is diamagnetic with no unpaired electrons.

64. 64 26 Fe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 Iron is paramagnetic with 4 unpaired electrons

65. 65 Details n Elements in the same column have the same electron configuration. n Elements in same groups have similar properties because of similar electron configuration. n Noble gases have filled energy levels. n Transition metals are filling the d orbitals.

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69. 69 1H1H 3 Li 11 Na 19 K 37 Rb 55 Cs 87 Fr 4 Be 12 Mg 20 Ca 38 Sr 56 Ba 88 Ra 41 Nb 21 Sc 39 Y 71 Lu 103 Lr 23 V 22 Ti 40 Zr 72 Hf 104 Rf 73 Ta 105 Db 24 Cr 42 Mo 74 W 106 Sg 25 Mn 43 Tc 75 Re 107 Bh 44 Ru 26 Fe 76 Os 108 Hs 27 Co 45 Rh 77 Ir 109 Mt 28 Ni 46 Pd 78 Pt 110 29 Cu 47 Ag 79 Au 111 30 Zn 48 Cd 80 Hg 112 13 Al 83 Bi 51 Sb 33 As 15 P 114 82 Pb 50 Sn 32 Ge 14 Si 10 Ne 9F9F 8O8O 7N7N 6C6C 86 Rn 81 Tl 49 In 31 Ga 5B5B 84 Po 116 85 At 54 Xe 36 Kr 53 I 35 Br 52 Te 34 Se 18 Ar 2 He 17 Cl 16 S 59 Pr 58 Ce 61 Pm 60 Nd 57 La 62 Sm 63 Eu 64 Gd 65 Tb 66 Dy 67 Ho 70 Yb 69 Tm 68 Er 91 Pa 90 Th 93 Np 92 U 89 Ac 94 Pu 95 Am 96 Cm 97 Bk 98 Cf 99 Es 102 No 101 Md 100 Fm s block d block p block f block