1. Modern Atomic Theory

2. Rutherford’s Atom Rutherford and his coworkers were able to show that the nucleus of the atom is composed of Protons and Neutrons. Rutherford also found that the nucleus is surrounded by a cloud of Electrons. What and where exactly are the electrons.

3. Electromagnetic Radiation Energy travels through space as waves OR as a stream of tiny packets of energy called photons.

4. Wave Types From Highest Energy to Lowest Gamma Rays X-Rays Ultra-Violet Visible Light Infrared Microwave Radio

5. Visible Light Violet – about 400 nm or 4.00 x 10 -7 m Indigo – about 445 nm or 4.45 x 10 -7 m Blue – about 475 nm or 4.75 x 10 -7 m Green – about 510 nm or 5.10 x 10 -7 m Yellow – about 570 nm or 5.70 x 10 -7 m Orange – about 590 nm or 5.90 x 10 -7 m Red – about 650 nm or 6.50 x 10 -7 m

6. Color Blindness Red-Green Blue-Yellow

7. Parts of a Wave - Wavelength wavelength  – distance between 2 wave peaks

8. Parts of a Wave – Frequency and Speed frequency  – how many wave peaks pass a certain point per given time period speed – how fast a given peak travels thru space In general, longer, lower

9. Emission of Energy by Atoms When atoms absorb energy from some source they become excited; They jump from a lower level to a higher level they can release this energy by emitting photons of light (energy of light = energy change of atom) They return to the lower level

10. Velocity (speed) Equation Velocity = Frequency x Wavelength V= f x Speed of light (c) = 3.0 x 10 8 m/s Frequency is calculated in Hertz (hz) or inverse seconds (1/s)

11. Practice Problems A light wave was determined to have a wavelength of 4.00 x 10 -7 m while traveling at the speed of light. What is the frequency of this wave?

12. Solution Wavelength ( ): 4.00 x 10 -7 m Velocity: 3.0 x 10 8 m/s Equation: V= f x f = V/ f = 3.0 x 10 8 m/s 4.00 x 10 -7 m f = 7.50 x 10 14 Hz

13. Wave Energy Equation The energy of the wave or light can also be calculated. Equation: Energy = Planck's constant x frequency e = h x f h = 6.626 x 10 -34 Joule*seconds (J*s)

14. Practice What is the Energy of a wave that has a frequency of 7.50 x 10 14 Hz?

15. Solution Equation: e = h * f e = x h = 6.62606896 x 10 -34 J*s f = 7.50 x 10 14 Hz E = 6.626 x 10 -34 J*s * 7.50 x 10 14 Hz E = 4.97 x 10 -19 J

16. Examples Example: flame tests: Na burns yellow; Li – red; Cu – green http://www.dvaction.org/mediaplayback.php?id=225

17. Another Problem What is the Energy in a photon of light from Sodium?

18. Solution What is the Equation? e = h * f What do we know? h (always 6.626 x 10 -34 J*s What do we need? Frequency (f)

19. Getting frequency How do we get frequency? f = V/ v = 3.0 x 10 8 m/s (it is light)  5.70 x 10 -7 m F = 3.0 x 10 8 m/s / 5.70 x 10 -7 m F = 5.26 x 10 14 Hz

20. Solving E = h x f E = 6.626 x 10 -34 J*s x 5.26 x 10 14 Hz E = 3.49 x 10 -19 J

21. Energy Levels of Hydrogen If we look at photons emitted after H absorbs energy, we see only certain types of photons which means only certain energy changes occur and H must have discrete energy levels. This means that only certain values are allowed – quantized, not continuous. Analogy: ramp – varies continuously in elevation; stairs – allow only certain elevations (quantized)

22. Light Spectrums Two Types: Continuous Spectrums: All color shown, like a rainbow Line Spectrums: Only certain colors shown as lines

23. Continuous All colors of light shown with no breaks

24. Line Spectrums Only certain colors shown, with black areas between

25. Atoms give off Line Spectrums What does this mean? Neils Bohr solved this in 1913 He said the spectrums showed that electrons jumped from level to level as they moved

26. The Bohr Model of the Atom Electrons have only specific Energy Levels (n) He said that electrons orbited like planets Could only have certain orbits

27. Bohr Model of the Atom Pictures the atom as a small positive nucleus with electrons orbiting around it; electrons can jump from orbit to orbit when they absorb or emit energy. This model only worked for hydrogen.

28. Bohr’s Explanation The atom looks like the solar system, and the electrons orbit around the nucleus like planets. The line spectrum shows the energy levels (orbits) that the electrons can be in (occupy). Bohr knew something wasn’t right with his explanation, but it was the best he could come up with.

29. Wave Mechanical Model of the Atom or Quantum Mechanics Schrodinger and de Broglie suggested that perhaps electrons, like light, could act as particles or waves; quantum mechanics is a math model that predicts the possibilities of finding the electron at given points in space around the nucleus

30. Hydrogen Orbitals Orbital or Energy Level (n) (replaces orbit in Bohr model) – volume of space around nucleus where the electron is most likely found. (90% probability) Now we say electrons absorb and emit energy when moving from orbital to orbital

31. Energy Levels and the Periodic Table The Orbital or Principle Quantum Number is called n. N depends on the row and block Blockn sRow Number p dN = Row - 1 fN = Row - 2

32. Blocks on the Periodic Table Exceptions: He is an s. All other group 18’s are p. La and Ac are d’s.

33. Examples Hydrogen: 1 st Row, s block. N = 1 Iron: 4 th Row, d block. N = 4 – 1 = 3 Bromine: 4 th Row, p block. N = 4 Gold: 6 th Row, d block. N = 6 – 1 = 5 Uranium: 7 th Row, f block. N = 7 – 2 = 5

34. Sublevels or Angular Momentum ( l ) Sublevels (second quantum number) ( l ) are the blocks on the table. They tell us the shape of the orbital, or how the electron moves. There are 4 types: s, p, d, f based on energy level

35. Sublevels Energy Level Sublevels 1s 2S, p 3S, p, d 4 +S, p, d, f

36. s S is shaped like a ball

37. p P is shaped like a dumb bell or figure 8

38. D and F D and F are shaped like a propeller

39. Second Quantum Numbers blockNumber s0 p1 d2 f3

40. Third Quantum Number (m l ) Magnetic Quantum Number (m l ) tells us the three dimensional orientation (x, y, z axis) M l is – l to l If l is m l can be 00 1-1, 0, 1 2 -2, -1, 0, 1, 2 3 -3,-2,-1,0,1,2

41. Fourth Quantum Number – Spin (m s ) Spin tells us whether the electron is spinning clockwise (up) or counter clockwise (down) down is -1/2, up is +1/2 The first half of any block is -1/2, the second half is +1/2

42. Example 2 nd Row, p block has 6 elements, the first 3 are -1/2, the last 3 are +1/2 4 th Row, d block has 10 elements, the first 5 are -1/2, the last 5 are +1/2

43. Hydrogen Orbitals (cont) Principal Quantum Number or energy levels or shells are labeled as 1, 2, 3, 4… and are divided into sublevels or subshells s,p,d,f 1. Shell 1 has one subshell – 1s (1 spherical orbital) 2. Shell 2 has 2 subshells – 2s (1 spherical orbital) and 2p (3 figure 8 shaped orbitals) p 291 picture 3. Shell 3 has 3 subshells – 3s (1 spherical orbital) and 3p (3 figure 8 orbitals) and 3d (5 orbitals) p 292 picture

44. Hydrogen Orbitals (cont) As the energy level or shell number increases, size of orbitals increases and electrons are more likely to be further from the nucleus If more than one electron is present, each orbital can hold a maximum of 2 electrons of opposite spin – Pauli Exclusion Principle

45. Electron Arrangements in Atoms Put electrons in lowest orbitals first (1s, 2s, 2p, 3s, 3p) Can show electron arrangements 3 ways: 1. Electron configuration 2. Orbital diagram 3. Abbreviated electron configuration (Nobel Gas Configuration)

46. Electron Configurations The four quantum numbers n = energy level l = orbital (shape) m l = sublevel (whether it is by itself or paired up) m s = spin (up or down)

47. Energy Levels 1 st Energy level holds 2 electrons 2 nd Energy level holds 8 electrons 3 rd Energy level holds 18 electrons 4 th Energy level hold 32 electrons 5 th Energy level holds 32 electrons 6 th Energy level holds 32 electrons 7 th Energy level holds 32 electrons

48. Energy and Orbitals

49. Energy levels Correspond to the row on the periodic table. Indicate how far from the nucleus the electron is.  i.e. – which orbit it is in

50. Orbital (sublevel) The shape of the orbit in which the electron moves S = sphere, like a baseball  The electron would be found somewhere on the surface of the baseball p, d, f – moves in a figure 8 pattern.

51. m l = Sublevel Each sublevel has a maximum amount of electrons it can hold Each s can hold 2 e - Each p can hold 6 e - Each d can hold 10 e - Each f can hold 14 e - Aufbau Principle - One sublevel doesn’t get any electrons till the one before fills up. The number of electrons is written as a superscript.

52. Order to fill based on increasing energy 1s 2s, 2p 3s, 3p 4s, 3d, 4p 5s, 4d, 5p 6s, 4f, 5d, 6p 7s, 5f, 4d, 7p Notice, the energy level for d, is one less then the row, and for f, it is two less then the row.

53. Filling Diagram for Sublevels Aufbau Principle

54. Examples H He Li C K Fe

55. m s = spin The direction the electron spins, up or down The first half of the electrons in a sublevel always spin up, the second half always spin down

56. Orbital Diagrams Pictures of the electron configuration using ___ for each sublevel, and arrows for spin Examples:  H  He  Li  C  K  Fe

57. Abbreviated Electron Configuration Also known as Nobel Gas Configurations write previous noble gas symbol in brackets to represent innermost electrons then the remainder of the electron configuration Example: O: [He]2s 2 2p 4

58. Valence vs Core Electrons valence electrons – electrons in the outermost shell that are involved in chemical reactions Example: O has 6 valence electrons (2 in 2s and 4 in 2p) core electrons – inner electrons not involved in bonding or chemical reactions Example: O has 2 core electrons in the 1s orbital

59. Position in Periodic Table First two columns filling s orbitals, next ten columns d orbitals, last six columns p orbitals and bottom two rows f orbitals. The shell number corresponds to the row number for columns 1, 2, 13-18. Column 3- 12 begin with the 3rd shell. The bottom two rows begin with the 4th shell. Note: Columns 1, 2, 13-18 are called the representative elements and have 1-8 electrons respectively

60. Periodic Properties - Trends Metallic and nonmetallic activity - most active metals (most likely to lose electrons and form + ions) are at lower left of Periodic Table Example: Ga > Al; Li > Be most active nonmetals (most likely to gain electrons and form – ions) at upper right of Table Example: Br > I; I > Te

61. More Trends Ionization energy – energy required to remove an electron (in the gas phase) – lower left, smallest ionization energy (metals want to lose electrons and further from the nucleus are easier to remove) – upper right, largest ionization energy example: Br > I Be > Li

62. More Atomic size (radius) lower left, largest size upper right, smallest size As you go down a group, the number of energy levels containing electrons increases, increasing size As you go across a period, you are adding electrons to same energy level but more protons are present to attract the electrons more strongly, making atomic size smaller Example: I > BrLi > Be