1. Ernest Rutherfords Model l Discovered dense positive piece at the center of the atom- nucleus l Electrons would surround and move around it, like planets around the sun l Atom is mostly empty space l It did not explain the chemical properties of the elements – a better description of the electron behavior was needed

2. Niels Bohrs Model l Why dont the electrons fall into the nucleus? l Move like planets around the sun. In specific circular paths, or orbits, at different levels. An amount of fixed energy separates one level from another.

3. The Bohr Model of the Atom Niels Bohr I pictured the electrons orbiting the nucleus much like planets orbiting the sun. However, electrons are found in specific circular paths around the nucleus, and can jump from one level to another.

4. Bohrs model l Energy level of an electron analogous to the rungs of a ladder l The electron cannot exist between energy levels, just like you cant stand between rungs on a ladder l A quantum of energy is the amount of energy required to move an electron from one energy level to another

5. The Quantum Mechanical Model l Energy is quantized - It comes in chunks. l A quantum is the amount of energy needed to move from one energy level to another. l Since the energy of an atom is never in between there must be a quantum leap in energy. l In 1926, Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom

6. Schrodingers Wave Equation probability Equation for the probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger

7. l Things that are very small behave differently from things big enough to see. l The quantum mechanical model is a mathematical solution l It is not like anything you can see (like plum pudding!) The Quantum Mechanical Model

8. l Has energy levels for electrons. l Orbits are not circular. l It can only tell us the probability of finding an electron a certain distance from the nucleus. The Quantum Mechanical Model

9. l The atom is found inside a blurry electron cloud l An area where there is a chance of finding an electron. l Think of fan blades The Quantum Mechanical Model

10. Atomic Orbitals l Principal Quantum Number (n) = the energy level of the electron: 1, 2, 3, etc. l Within each energy level, the complex math of Schrodingers equation describes several shapes. l These are called atomic orbitals (coined by scientists in 1932) - regions where there is a high probability of finding an electron. l Sublevels- like theater seats arranged in sections: letters s, p, d, and f

11. Principal Quantum Number Generally symbolized by n, it denotes the shell (energy level) in which the electron is located. Maximum number of electrons that can fit in an energy level is: 2n 2 How many e - in level 2? 3?

12. Summary s p d f # of shapes (orbitals) Maximum electrons Starts at energy level 1 2 1 3 6 2 5 10 3 7 14 4

13. By Energy Level l First Energy Level l Has only s orbital l only 2 electrons l 1s 2 l Second Energy Level l Has s and p orbitals available l 2 in s, 6 in p l 2s 2 2p 6 l 8 total electrons

14. By Energy Level l Third energy level l Has s, p, and d orbitals l 2 in s, 6 in p, and 10 in d l 3s 2 3p 6 3d 10 l 18 total electrons l Fourth energy level l Has s, p, d, and f orbitals l 2 in s, 6 in p, 10 in d, and 14 in f l 4s 2 4p 6 4d 10 4f 14 l 32 total electrons

15. By Energy Level Any more than the fourth and not all the orbitals will fill up. l You simply run out of electrons l The orbitals do not fill up in a neat order. l The energy levels overlap l Lowest energy fill first.

16. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f aufbau diagram - page 133 Aufbau is German for building up

17. Electron Configurations… l …are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how: 1) Aufbau principle - electrons enter the lowest energy first. This causes difficulties because of the overlap of orbitals of different energies – follow the diagram! 2) Pauli Exclusion Principle - at most 2 electrons per orbital - different spins

18. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli To show the different direction of spin, a pair in the same orbital is written as:

19. Electron Configurations 3) Hunds Rule- When electrons occupy orbitals of equal energy, they dont pair up until they have to. l Lets write the electron configuration for Phosphorus We need to account for all 15 electrons in phosphorus

20. l The first two electrons go into the 1s orbital Notice the opposite direction of the spins l only 13 more to go... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

21. l The next electrons go into the 2s orbital l only 11 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

22. The next electrons go into the 2p orbital only 5 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

23. The next electrons go into the 3s orbital only 3 more... Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f

24. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f The last three electrons go into the 3p orbitals. They each go into separate shapes (Hunds) 3 unpaired electrons = 1s 2 2s 2 2p 6 3s 2 3p 3 Orbital notation

25. Section 5.3 Physics and the Quantum Mechanical Model l OBJECTIVES: Identify the source of atomic emission spectra.

26. Section 5.3 Physics and the Quantum Mechanical Model l OBJECTIVES: Distinguish between quantum mechanics and classical mechanics.

27. Light l The study of light led to the development of the quantum mechanical model. l Light is a kind of electromagnetic radiation. l Electromagnetic radiation includes many types: gamma rays, x-rays, radio waves… l Speed of light = 2.998 x 10 8 m/s, and is abbreviated c l All electromagnetic radiation travels at this same rate when measured in a vacuum

28. - Page 139 R O Y G B I V Frequency Increases Wavelength Longer

29. Parts of a wave Wavelength Amplitude Origin Crest Trough

30. Equation: c = c = speed of light, a constant (2.998 x 10 8 m/s) (nu) = frequency, in units of hertz (hz or sec -1 ) (lambda) = wavelength, in meters Electromagnetic radiation propagates through space as a wave moving at the speed of light.

31. Wavelength and Frequency l Are inversely related As one goes up the other goes down. l Different frequencies of light are different colors of light. l There is a wide variety of frequencies l The whole range is called a spectrum

32. Radio waves Micro waves Infrared. Ultra- violet X- Rays Gamma Rays Low Frequency High Frequency Long Wavelength Short Wavelength Visible Light Low Energy High Energy

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

34. Atomic Spectra l White light is made up of all the colors of the visible spectrum. l Passing it through a prism separates it.

35. If the light is not white l By heating a gas with electricity we can get it to give off colors. l Passing this light through a prism does something different.

36. Atomic Spectrum l Each element gives off its own characteristic colors. l Can be used to identify the atom. l This is how we know what stars are made of.

37. These are called the atomic emission spectrum Unique to each element, like fingerprints! Very useful for identifying elements

38. Light is a Particle? l Energy is quantized. l Light is a form of energy. l Therefore, light must be quantized l These smallest pieces of light are called photons. l Photoelectric effect? Albert Einstein l Energy & frequency: directly related.

39. Equation: E = h E = Energy, in units of Joules (kg·m 2 /s 2 ) (Joule is the metric unit of energy) (Joule is the metric unit of energy) h = Plancks 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.

40. Explanation of atomic spectra l When we write electron configurations, we are writing the lowest energy. l The energy level, and where the electron starts from, is called its ground state - the lowest energy level.

41. Changing the energy l Lets look at a hydrogen atom, with only one electron, and in the first energy level.

42. Changing the energy l Heat, electricity, or light can move the electron up to different energy levels. The electron is now said to be excited

43. Changing the energy l As the electron falls back to the ground state, it gives the energy back as light

44. Experiment #6, page 49-

45. l They may fall down in specific steps l Each step has a different energy Changing the energy

46. { { {

47. l The further they fall, more energy is released and the higher the frequency. l This is a simplified explanation! l The orbitals also have different energies inside energy levels l All the electrons can move around. Ultraviolet Visible Infrared

48. What is light? l Light is a particle - it comes in chunks. l Light is a wave - we can measure its wavelength and it behaves as a wave If we combine E=mc 2, c=, E = 1/2 mv 2 and E = h, then we can get: = h/mv (from Louis de Broglie) l called de Broglies equation l Calculates the wavelength of a particle.

49. Wave-Particle Duality J.J. 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!

50. Confused? Youve Got Company! No familiar conceptions can be woven around the electron; something unknown is doing we dont know what. Physicist Sir Arthur Eddington The Nature of the Physical World 1934

51. The physics of the very small l Quantum mechanics explains how very small particles behave Quantum mechanics is an explanation for subatomic particles and atoms as waves l Classical mechanics describes the motions of bodies much larger than atoms

52. Heisenberg Uncertainty Principle l It is impossible to know exactly the location and velocity of a particle. l The better we know one, the less we know the other. l Measuring changes the properties. l True in quantum mechanics, but not classical mechanics

53. 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

54. It is more obvious with the very small objects l To measure where a electron is, we use light. l But the light energy moves the electron l And hitting the electron changes the frequency of the light.

55. Moving Electron Photon Before Electron velocity changes Photon wavelength changes After Fig. 5.16, p. 145