1. Chapter 12 Electrons in Atoms

2. Greek Idea lDlDemocritus and Leucippus l Matter is made up of indivisible particles lDlDalton - one type of atom for each element

4. Thomson’s Model l Discovered electrons l Atoms were made of positive stuff l Negative electron floating around l “Plum-Pudding” model

6. Rutherford’s Model l Discovered dense positive piece at the center of the atom l Nucleus lElElectrons moved around l Mostly empty space

8. Bohr’s Model l Why don’t the electrons fall into the nucleus? l Move like planets around the sun. l In circular orbits at different levels. l Amounts of energy separate one level from another.

9. Bohr’s Model Nucleus Electron Orbit Energy Levels

10. Bohr’s Model Increasing energy Nucleus First Second Third Fourth Fifth } l Further away from the nucleus means more energy. l There is no “in between” energy l Energy Levels

12. The Quantum Mechanical Model l Energy is quantized. It comes in chunks. l A quanta 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 Schrödinger derived an equation that described the energy and position of the electrons in an atom

13. 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. The Quantum Mechanical Model

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

15. l The atom is found inside a blurry “electron cloud” l A area where there is a chance of finding an electron. l Draw a line at 90 % The Quantum Mechanical Model

18. Atomic Orbitals l Principal Quantum Number (n) = the energy level of the electron. l Within each energy level the complex math of Schrödinger's equation describes several shapes. l These are called atomic orbitals l Regions where there is a high probability of finding an electron.

19. l There is 1 s orbital for every energy level l Spherical shaped l Each s orbital can hold 2 electrons. l Called the 1s, 2s, 3s, etc.. orbitals. S orbitals

20. P orbitals l Start at the second energy level l 3 different directions l 3 different shapes l Each can hold 2 electrons

22. P Orbitals

23. D orbitals l Start at the third energy level l 5 different shapes l Each can hold 2 electrons

26. F orbitals l Start at the fourth energy level l Have seven different shapes l 2 electrons per shape

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

29. By Energy Level l Third energy level l 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 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

30. By Energy Level l 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.

31. 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Increasing Energy

32. Electron Configurations l The way electrons are arranged in atoms. l Aufbau principle- electrons enter the lowest energy first. l This causes difficulties because of the overlap of orbitals of different energies. l Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

33. Electron Configuration l Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to. l Let’s determine the electron configuration for Phosphorus l Need to account for 15 electrons

34. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4 p 4 What is it? Selenium

35. Orbitals fill in order l Lowest energy to higher energy. l Adding electrons can change the energy of the orbital. l Half filled orbitals have a lower energy. l Makes them more stable. l Changes the filling order

36. Write these electron configurations l Titanium - 22 electrons l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 Vanadium - 23 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 l Chromium - 24 electrons l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 is expected l But this is wrong!! Why??

37. 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 kinds of waves l All move at 3.00 x 10 8 m/s = C

38. Parts of a wave Wavelength Amplitude Origin Crest Trough

39. Parts of Wave l Origin - the base line of the energy. l Crest - high point on a wave l Trough - Low point on a wave l Amplitude - distance from origin to crest l Wavelength - distance from crest to crest l Wavelength - is abbreviated  Greek letter lambda.

40. Frequency l The number of waves that pass a given point per second. l Units are cycles/sec or hertz (hz) l Abbreviated  the Greek letter nu c =

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

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

43. Atomic Spectrum How color tells us about atoms

44. Prism l White light is made up of all the colors of the visible spectrum. l Passing it through a prism separates it.

45. 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.

46. 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!

47. Confused??? You’ve Got Company! “No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.” Physicist Sir Arthur Eddington The Nature of the Physical World 1934

48. 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.

49. 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.

50. Types of electromagnetic radiation:

51. 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.

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

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

55. …produces a “bright line” spectrum Spectroscopic analysis of the hydrogen spectrum…

56. This produces bands of light with definite wavelengths. Electron transitions involve jumps of definite amounts of energy.