1. Chapter 5 “Electrons in Atoms”

2. 5.1 Models of the Atom l OBJECTIVES: Identify the inadequacies in the Rutherford atomic model. Identify the new proposal in the Bohr model of the atom. Describe the energies and positions of electrons according to the quantum mechanical model. Describe the shapes of orbitals related to different sublevels differ.

3. Ernest Rutherford’s 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

4. Niels Bohr’s Model l Why don’t electrons fall into the nucleus (opposites charges attract, right?) 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.  The greater the energy of the electron, the farther it would be from the nucleus

5. 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.

6. The emission spectrum of hydrogen is composed of only a few wavelengths of light. Each color corresponds to a precise amount of energy. Bohr correctly connected the production of the specific colors to electrons falling from a higher to a lower energy level Emission spectrum of H

7. Bohr’s model l Energy level of an electron analogous to the rungs of a ladder l Electrons cannot exist between energy levels, just like you can’t stand between rungs on a ladder Electron energy levels in atoms are like the rungs on the ladder in the picture. The closer the electron to the nucleus, the bigger the “step” to the next energy level.

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

9. 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

10. l Electrons occupy different energy levels within the electron cloud. l Orbits are not circular. l We can only tell the probability of finding an electron in a certain area around the nucleus. Quantum Mechanical Model The nucleus is found inside a blurry “electron cloud”

11. Atomic Orbitals l Within each energy level (principal quantum number), complex mathematical equations describe the regions where electrons are most likely to be found. l Atomic orbitals are the geographic areas around the nucleus where electrons of a particular energy level are most likely to be found. l Sublevels - like theater seats arranged in sections: letters s, p, d, and f. These “sections” are the areas where the electrons are “seated”

12. Electron Orbitals

13. Describing an Electron’s Position Principal Quantum Number (n) denotes the shell (energy level) in which the electron is located. The lower the energy level, the closer the electrons to the nucleus Maximum number of electrons that can fit in an energy level is: 2n 2 How many e - possible in level 2? 3?

14. By Energy Level l First Energy Level l Has only an s orbital l only 2 electrons (the max for any orbital) l If filled, it is abbreviated: 1s 2 l 2 nd Energy Level l Has one s and three p orbitals available l 2 e - in s, 6 in p l If second level filled = 2s 2 2p 6 l = 8 total electrons

15. By Energy Level l Third energy level l Has s, p, and d (5) 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

16. 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.

17. 5.2: Electron Configurations l OBJECTIVES: Use the periodic table to help you write the electron configuration for an atom/ion Identify atoms that have electron configurations that do not conform to the usual “rules”

18. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f Aufbau Diagram Aufbau is German for “building up”

19. 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

20. Electron Configurations 3) Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until they have to. l Let’s write the electron configuration for Phosphorus  We need to account for all 15 electrons in phosphorus

21. 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

22. 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

23. 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

24. 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

25. 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 (Hund’s) 3 unpaired electrons = 1s 2 2s 2 2p 6 3s 2 3p 3 Orbital notation

26. l An animation showing how electrons fill the orbitals of the first 48 elements can be found at: http://chemmovies.unl.edu/ChemAnime/ECONFIG/ECONFIG.html Electron Configurations (Just click on the above link)

27. Using the Periodic Table to Write Electron Configurations The electron configuration of Si ends with 3s 2 3p 2 The electron con- figuration of Rh ends with 5s 2 4d 7 Rh Si O Br Ti Write electron configurations for O, Ti, and Br.

28. Write the electron configurations for these elements: lTlTitanium - 22 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 lVlVanadium - 23 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 lClChromium - 24 electrons 11s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 (expected) BBut this is not the correct config!!

29. Chromium is actually: l 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 l Why? l This gives two sets of half filled orbitals l In this case, the half full orbitals are slightly lower in energy, so more stable. l Similarly, Cu and Ag are more stable with completely filled d-sublevels  Copper: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10  Silver: [Kr] 5s 1 4d 10

30. Irregular configurations of Cr and Cu Chromium “steals” a 4s electron to make its 3d sublevel HALF FULL (each orbital with one electron) Copper “steals” a 4s electron to FILL its 3d sublevel

31. 5.3 Quantum Mechanical Model l OBJECTIVES: Describe the relationship between the wavelength and frequency of light. Identify the source of atomic emission spectra. Explain how the frequencies of emitted light are related to changes in electron energies. Distinguish between quantum mechanics and classical mechanics.

32. Flame Tests l When different elements are heated in a Bunsen flame, they give off different colors of light. Why??? Cu Na K

33. 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

34. - Page 139 “R O Y G B I V” Frequency and Energy increases Wavelength Longer, energy lower

35. Parts of a wave Wavelength Amplitude Origin Crest Trough

36. 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.

37. Wavelength and Frequency l Are inversely related As one goes up, 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

38. 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

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

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

41. 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.

42. Atomic Emission Spectra 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.

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

44. Light is a Particle? l Energy is quantized. There is a smallest “piece” of energy = quantum l Light is a form of energy, therefore, light must be quantized l The smallest “pieces” of light are called photons. l Photoelectric effect? Einstein helped show the particle nature of light

45. Photoelectric Effect l Red wavelength of light does not have enough energy to eject an electron l Green wavelength has enough energy to eject electrons from the metal l Photon of violet light, with even higher energy, ejects an electron at a higher speed

46. 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 = 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.

47. The Math in Chapter 5 l There are 2 equations: 1) c = 2) E = h Know these!

48. Examples 1) What is the wavelength of blue light with a frequency of 8.3 x 10 15 hz? 2) What is the frequency of red light with a wavelength of 4.2 x 10 -5 m? 3) What is the energy of a photon of each of the above?

49. 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 it’s ground state - the lowest energy level.

50. Changing the energy l Let’s look at a hydrogen atom, with only one electron, and in the first energy level.

51. 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 ”

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

53. Why does hamburger have lower energy than steak? Because it's in the ground state.

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

55. Confused? You’ve Got Company! I do not like it, and I am sorry I ever had anything to do with it. Erwin Schrödinger (1887-1961) Nobel Prize, 1933. Speaking of quantum mechanics

56. 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

57. 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

58. 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

59. 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.

60. Moving Electron Photon Before Electron velocity changes Photon wavelength changes After The atomic world is a weird place!