2. Chemistry Chapter 4 Arrangement of Electrons in AtomsArrangement of Electrons in Atoms The 1998 Nobel Prize in Physics was awarded "for the discovery of a new form of quantum fluid with fractionally charged excitations." At the left is a computer graphic of this kind of state.1998 Nobel Prize in Physics

3. Pop Quiz 1.List one part of Dalton’s Atomic Theory that has been disproven and explain the Modern Atomic Theory for that part. 2.Thompson and Millikan’s experiments led to conclusions about the electron. List one conclusion and why they made that conclusion. 3.How many protons, neutrons, and electrons does a neutral atom of Hydrogen-3 have?

4. The Puzzle of the Atom The Rutherford model of the atom was an improvement over previous models, but it was incomplete. Protons and electrons are attracted to each other because of opposite charges Atoms have a positive center and negative surrounding Despite these facts, atoms don’t collapse. WHY??????

5. In the early 1900’s, a new atomic model evolved as a result of investigations into the absorption and emission of light by matter. The study revealed an intimate relationship between light and atom’s electrons and began to explain how atom’s are organized. Visible light –A kind of electromagnetic radiation. Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space. A wave is a repetitive disturbance that propagates through space.

6. Electromagnetic Waves Click in this box to enter notes. Copyright © Houghton Mifflin Company. All rights reserved. Go to Slide Show View (press F5) to play the video or animation. (To exit, press Esc.) This media requires PowerPoint® 2000 (or newer) and the Macromedia Flash Player (7 or higher). [To delete this message, click inside the box, click the border of the box, and then press delete.]

9. Amplitude is the distance from the baseline to a trough or a crest.

10. Wavelength ( ) is the distance between corresponding points on adjacent waves.

11. Frequency ( ) is defined as the number of waves that pass a given point in a specific time, usually one second.

12. Electromagnetic radiation propagates through space as a wave moving at the speed of light. Therefore frequency, wavelength and the speed of light are mathematically related. c = C = speed of light, a constant (3.00 x 10 8 m/s) = frequency, in units of hertz (hz, sec -1 ) = wavelength, in meters http://www.colorado.edu/physics/2000/waves_particles/lightspeed-1.htmlwww.colorado.edu/physics/2000/waves_particles/lightspeed-1.html

13. Practice Problem Determine the frequency of light with a wavelength of 4.257 × 10 -7 cm.

14. Prism analysis of the visible spectrum p Prism analysis of the visible spectrum produces all of the colors in a continuous spectrum Visible Light is a form a electromagnetic radiation. Between 400 nanometers to 700 nanometers

15. Refraction of White Light Click in this box to enter notes. Copyright © Houghton Mifflin Company. All rights reserved. Go to Slide Show View (press F5) to play the video or animation. (To exit, press Esc.) This media requires PowerPoint® 2000 (or newer) and the Macromedia Flash Player (7 or higher). [To delete this message, click inside the box, click the border of the box, and then press delete.]

16. Types of electromagnetic radiation: The Electromagnetic Spectrum. http://www.colorado.edu/physics/2000/waves_particles/index.htmlwww.colorado.edu/physics/2000/waves_particles/index.html

17. The Photoelectric Effect Observed in the early 1900’s and explained light as a particle. Refers to the emission of electrons from a metal when light shines on the metal. –For a given metal, no electrons were emitted if the light’s frequency was below a certain minimum.

18. Photoelectric Effect Click in this box to enter notes. Copyright © Houghton Mifflin Company. All rights reserved. Go to Slide Show View (press F5) to play the video or animation. (To exit, press Esc.) This media requires PowerPoint® 2000 (or newer) and the Macromedia Flash Player (7 or higher). [To delete this message, click inside the box, click the border of the box, and then press delete.]

19. The Photoelectric Effect Max Planck proposed that a hot object does not emit electromagnetic energy continuously, but in packets called quanta. –Quanta is the minimum quantity of energy that can be lost or gained by an atom. –Later Albert Einstein called these particles photons. A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy.

20. E = h E = Energy, in units of Joules (kg·m 2 /s 2 ) h = Plank’s constant (6.626 x 10 -34 J·s) = frequency, in units of hertz (hz, sec -1 ) Through experiment Plank determined that the energy (E ) of electromagnetic radiation is directly proportional to the frequency ( ) of the radiation.

21. Quick Review Wavelength is inversely proportional to frequency. Frequency is directly proportional to Energy.

22. Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table

23. Practice Problem Determine the energy in joules of a photon whose frequency is 3.55 × 10 17 Hz.

24. The Hydrogen-Atom Line- Emission Spectrum When current is passed through gas at low pressure, the potential energy of some of the gas atoms increases. Atoms increased from the ground state (the lowest energy state level of an atom) to an excited state (a state in which the atom has a higher potential energy than the ground state).

25. The Hydrogen-Atom Line- Emission Spectrum When an excited atom returns to its ground state, it gives off the energy it gained in the form of light. –Neon Signs.

26. …produces a line-emission spectrum Prism analysis of the hydrogen spectrum…

28. …produces a line-emission spectrum Prism analysis of the hydrogen spectrum…

29. Bohr Model of the Hydrogen Atom http://www.colorado.edu/physics/2000/quantumzone/lines2.htmlwww.colorado.edu/physics/2000/quantumzone/lines2.html

31. Pop Quiz 1.Isotopes of the same element have the same number of _________ and different numbers of ________. 2.If a wave has a long wavelength, how would you describe the frequency and energy? 3.How many protons, neutrons, and electrons does a neutral atom of Hydrogen-1 have? 4.Describe either Heisenberg Uncertainty Principle or Pauli Exclusion Principle.

32. The Bohr Model of the Atom Neils Bohr I pictured electrons orbiting the nucleus much like planets orbiting the sun. But I was wrong! They’re more like bees around a hive. WRONG!!!

33. Heisenberg Uncertainty Principle You can find out where the electron is, but not where it is going. OR… You can find out where the electron is going, but not where it is! “One cannot simultaneously determine both the position and momentum of an electron.” Werner Heisenberg

34. Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it.  Principal quantum number  Angular momentum quantum number  Magnetic quantum number  Spin quantum number

35. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli

36. Principal Quantum Number Generally symbolized by n, it denotes the shell (energy level) in which the electron is located. Number of electrons that can fit in a shell: 2n 2 Energy level 2 = 2(2) 2 = 8 electrons

37. Angular Momentum Quantum Number The angular momentum quantum number, generally symbolized by l, denotes the orbital shape (subshell) in which the electron is located. l Letter 0 s 1 p 2 d 3 f

38. Schrodinger Wave Equation probability Equation for probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger http://www.colorado.edu/physics/2000/quantumzone/schroedinger.htmlwww.colorado.edu/physics/2000/quantumzone/schroedinger.html

39. Orbital shapes are defined as the surface that contains 90% of the total electron probability. An orbital is a region within an atom where there is a probability of finding an electron.

40. Magnetic Quantum Number The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.

41. s orbital shape

42. P orbital shape

43. d orbital shapes

44. Shape of f orbitals

45. Assigning the Numbers  The three quantum numbers (n, l, and m) are integers.  The principal quantum number (n) cannot be zero.  n must be 1, 2, 3, etc.  The angular momentum quantum number (l) can be any integer between 0 and n - 1.  For n = 3, l can be either 0, 1, or 2.  The magnetic quantum number (m) can be any integer between -l and +l.  For l = 2, m can be either -2, -1, 0, +1, or +2.

46. Principle, angular momentum, and magnetic quantum numbers: n, l, and m l

47. Spin Quantum Number Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field. Possibilities for electron spin:

49. Electron Configurations The arrangement of electrons in an atom. Since atoms of different elements have different numbers of electrons, a distinct electron configuration exists for the atoms of each element.

50. Electron Configurations Three Rules for adding electrons to orbitals: –Aufbau principle – an electron occupies the lowest-energy orbital available. –Pauli exclusion principle – no two electrons in the same atom can have the same set of four quantum numbers. –Hund’s rule- One electron enters each orbital of equal energy until all orbitals contain one electron & all electrons in singly occupied orbitals must have the same spin.

51. Electron Configurations There are 3 ways to write electron configurations: –Orbital Notation –Electron-Configuration Notation –Nobel-Gas Notation

52. Aufbau principle an electron occupies the lowest- energy orbital available.

53. Orbital filling table

54. Pauli exclusion principle no two electrons in the same atom can have the same set of four quantum numbers Orbital Electron or He 1s Which orbital? Orbital Notation

55. Electron-Configuration Notation Eliminates the lines and arrows. The number of e - in a sublevel is shown by adding a superscript to the sublevel designation. So for He…… He 1s We rewrite He as 1s 2 where 1 represents the energy level, s represents the orbital shape and 2 represents the number of electrons 1s 2

56. Hund’s Rule One electron enters each orbital of equal energy until all orbitals contain one electron & all electrons in singly occupied orbitals must have the same spin. N 1s2s2p Orbital Notation Electron-Configuration Notation:1s2s2p 223

57. Practice Si Be P 1s 2s 2p3s3p1s 2s 1s 2s 2p3s3p 1s 3p2s 2p 3s 1s 22262 2s 22 1s 3p2s3s 22263 2p

58. Practice Kr 1s 2s 2p3s4p 4s 3d 3p 1s 3p2s3s 22266 2p 4s3d 4p 2 106

59. Irregular confirmations of Cr and Cu Chromium steals a 4s electron to half fill its 3d sublevel Copper steals a 4s electron to FILL its 3d sublevel

60. 1s 2s 2p3s 4s 3d 3p Cr

61. Cu 1s 2s 2p3s 4s 3d 3p

62. Pop Quiz Please write the orbital notation and electron configuration for Gold. Define the four quantum numbers and describe how to find the first three.

63. Nobel-Gas Notation Simplifies electron-configuration notation For example: P 1s 3p2s3s 22263 2p P Ne 1s 2s 226 2p [Ne]3s 2 3p 3