2. The Puzzle of the Atom  Protons and electrons are attracted to each other because of opposite charges  Electrically charged particles moving in a curved path give off energy  Despite these facts, atoms don’t collapse

3. Wave-Particle Duality JJ Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave!

4. The Wave-like Electron Louis deBroglie The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.

5. 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 Electromagnetic radiation propagates through space as a wave moving at the speed of light.

6. Electromagnetic radiation.

7. Electromagnetic Radiation Most subatomic particles behave as PARTICLES and obey the physics of waves.Most subatomic particles behave as PARTICLES and obey the physics of waves.

8. Waves have a frequencyWaves have a frequency Use the Greek letter “nu”,, for frequency, and units are “cycles per sec”Use the Greek letter “nu”,, for frequency, and units are “cycles per sec” All radiation: = c where c = velocity of light = 3.00 x 10 8 m/secAll radiation: = c where c = velocity of light = 3.00 x 10 8 m/sec Electromagnetic Radiation

9. Types of electromagnetic radiation:

10. The Electromagnetic Spectrum AM radio Short wave radio Television channels FM radio Radar Microwave Radio Waves Gamma Rays X- Raysinfrared Increasing photon energy Increasing frequency Decreasing wavelength Red Orange Yellow Green Blue Indigo Violet UV Rays VisibleLightVisibleLight R O Y G B I V

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

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

13. Relating Frequency, Wavelength and Energy Common re-arrangements:

14. Modern model of the atom Every element has a unique chemical behavior: carbon acts like carbon; iron like iron. In 1913, the Danish physicist Niels Bohr proposed a model of the atom that explained these differences, as well as other observations known at the time.

15. Line Emission Spectra of Excited Atoms Excited atoms emit light of only certain wavelengths The wavelengths of emitted light depend on the element.

16. …produces all of the colors in a continuous spectrum Spectroscopic analysis of the visible spectrum…

17. Spectrum of Excited Hydrogen Gas

18. Line Spectra of Other Elements

19. What did Bohr’s model propose? Bohr's central insight was that the electrons in an atom can have only certain, specific energies: Their energy is quantized. The higher an electron's energy, the farther the electron is from the nucleus. For this reason, he called the energy levels shells.

20. 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!!!

21. Physics and Chemistry Quantum Physics of atomic and subatomic particles Objects can have only certain energies Classical : objects can have any energy Physics

22. Quantum Mechanical Model of the Atom Mathematical laws can identify the regions outside of the nucleus where electrons are most likely to be found. These laws are beyond the scope of this class…

23. The closer the electron is to the nucleus the lower the electron’s energy and the more stable the atom is.

24. Bohr’s shells He assigned principal quantum numbers to each shell (n) n=1 for the first n=2 for the second

25. Electron Energy Level (Shell) Generally symbolized by n, it denotes the probable distance of the electron from the nucleus. Number of electrons that can fit in a shell: 2n 2

26. Maximum number of electrons The maximum number of electrons in each shell is 2n 2 Electrons in n=1 have the lowest energy and are the most stable

27. Main features of Bohr’s Model 1.Orbits get larger as the orbit number increases(the radius increases) 2.Electrons in n=1 orbit are the most stable with lowest energy 3.Each orbit hold a maximum of 2n 2

28. Main features of Bohr’s Model 4. Orbits are filled with electrons from the innermost shell on out. Electrons are added to each orbit until filled to capacity before moving to next level.

29. Periodicity and line spectra

31. What is wrong with Bohr’s model?

32. Subshells and electron configuration Subshells are closely spaced in energy and size Letters are assigned to the subshells(s, p, d, f) n=2 has s and p subshells Each shell has a different maximum amount of electron capacity

33. Subshells and electron configuration Not every level contains all the subshells n=1 contains only subshell “s” Every level gets an additional subshell

34. Maximum number of electrons s=2 p=6 d=10 f=14 Energy increases from s  p  d  f Maximum electron capacity for a shell is unchanged from Bohr’s original model

35. Subshells and electron configuration

36. 4S 3S 2S 1S 3P 2P

39. Electron configuration: exact way that the electrons are arranged in subshell. In the “d” block, n value for d subshell is 1 less than period number. In the “f” block, n value for f subshell is 2 less than period number.

41. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

42. Partial electronic configurations for K through Kr

43. “f” block

44. Electron configuration for Gadolinium From 1 st row: 1s 2 From 2 nd row: 2s 2 2p 6 From 3 rd row: 3s 2 3p 6 From 4 th row: 4s 2 3d 10 4p 6 From 5 th row: 5s 2 4d 10 5p 6 From 6 th row: 6s 2 5d 1 4f 7 Gd is 1 deep in d block and 7 in f block

45. Diagonal Rules s 3p 3d s 2p s 4p 4d 4f s 5p 5d 5f 5g? s 6p 6d 6f 6g? 6h? s 7p 7d 7f 7g? 7h? 7i? 1234567 Steps: 1.Write the energy levels top to bottom. 2.Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level. 3.Draw diagonal lines from the top right to the bottom left. 4.To get the correct order, follow the arrows! By this point, we are past the current periodic table so we can stop.

46. neon's electron configuration (1s 2 2s 2 2p 6 ) Shorthand Configuration [Ne] 3s 1 third energy level one electron in the s orbital orbital shape Na = [1s 2 2s 2 2p 6 ] 3s 1 electron configuration A B C D

47. Octet Rule Elements react to form compounds in such a way as to put 8 electrons in their outermost(valence)shell, giving them a valence electron configuration identical to a noble gas and making them exceptionally stable

48. Metal: element that tends to lose its valence electrons in chemical reactions (cation) Nonmetal: element that tends to gain valence electrons in a reaction(anion)

50. Quantum Physics  uncertainty principle Werner Heisenberg(1930s) It is impossible to know exactly where the electron is or even where it is going

51. Orbitals An orbital is a region within an atom where there is a probability of finding an electron.

52. Orbitals of the same shape (s, for instance) grow larger as n increases… Nodes are regions of low probability within an orbital. Sizes of s orbitals

53. The s orbital has a spherical shape centered around the origin of the three axes in space. s orbital shape

54. There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space. P orbital shape

55. Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbbells ” …and a “dumbbell with a donut”! d orbital shapes

56. Shape of f orbitals