1. Development of Atomic Models

2. Democritus Greek philosopher Greek philosopher 400 BC 400 BC

3. Question Is there a limit to the number of times matter could be divided? Is there a limit to the number of times matter could be divided?

4. Democritus Theory Eventually, you would reach a piece that was “indivisible” Eventually, you would reach a piece that was “indivisible” Named this smallest piece of matter “atomos,” meaning “not to be cut.” Named this smallest piece of matter “atomos,” meaning “not to be cut.”

5. Atomos  Small, hard particles.  Differ in shape and size for each substance

6. Aristotle and Plato All matter made up of combination of earth, fire, air and water. Aristotle

7. The Four Elements?? This concept influenced early chemists called alchemists. This concept influenced early chemists called alchemists.

8. Buried in History “Atomos” theory was ignored and forgotten for more than 2000 years! “Atomos” theory was ignored and forgotten for more than 2000 years!

9. John Dalton (early 1800’s) Performed careful scientific experiments. Performed careful scientific experiments. Coined the term “atom”. Coined the term “atom”.

10. Dalton’s Atomic Theory Matter is made of tiny indivisible particles called atoms. Matter is made of tiny indivisible particles called atoms. Atoms of an element are alike, and different from atoms of other elements. Atoms of an element are alike, and different from atoms of other elements.

11. Dalton’s Atomic Theory Compounds are atoms of different elements combined in fixed proportions. Compounds are atoms of different elements combined in fixed proportions. Chemical reactions involve rearrangement of atoms. Chemical reactions involve rearrangement of atoms. Atoms cannot be created or destroyed, but are conserved. Atoms cannot be created or destroyed, but are conserved.

12. Pages from Dalton’s Journal

13. Hard Spheres Dalton’s model is called the “Hard Spheres Model”

14. JJ Thomson (1897)

15. Thomson’ Experiments Studied “cathode rays” (electric current) in a “Crooke’s Tube”. Studied “cathode rays” (electric current) in a “Crooke’s Tube”. Fluorescent screen, shows how ray behaved in a magnetic field. Fluorescent screen, shows how ray behaved in a magnetic field.

16. Cathode Rays were negatively charged

17. Cathode Rays were particles

18. http://youtu.be/XU8nMKkzbT8 http://youtu.be/XU8nMKkzbT8 http://youtu.be/XU8nMKkzbT8 http://youtu.be/Z61zCaAFky4 http://youtu.be/Z61zCaAFky4 http://youtu.be/Z61zCaAFky4 http://youtu.be/IdTxGJjA4Jw http://youtu.be/IdTxGJjA4Jw http://youtu.be/IdTxGJjA4Jw

19. JJ is Awesome Concluded the negative “cathode ray” particles came from within atoms. Discovered the first subatomic particle (electron). Discovered the first subatomic particle (electron).

20. What about the Positive? But…matter is neutral. But…matter is neutral. Must be a positive charge in the atom to balance the negative. Must be a positive charge in the atom to balance the negative.

21. Plum Pudding Model Positively charged sphere with with negatively charged particles scattered throughout. Positively charged sphere with with negatively charged particles scattered throughout.

22. Yummy…

23. Ernest Rutherford (1908) Physicist who worked with the new field of radioactive emissions. Physicist who worked with the new field of radioactive emissions.

24. Different Types of Radiation Used a magnetic field to determine there were three types of radiation. Used a magnetic field to determine there were three types of radiation. Alpha (α) Alpha (α) Beta (β) Beta (β) Gamma (γ) Gamma (γ)

25. Charges of Radiation The radiation had different charges. The radiation had different charges. Identify the charge each type of radiation has.

26. Shot alpha particles, at a very thin piece of gold foil. Shot alpha particles, at a very thin piece of gold foil. These particles have a positive charge These particles have a positive charge Fluorescent screen shows where the particles went. Fluorescent screen shows where the particles went. Gold Foil Experiment

27. Observation: Observation: Almost all alpha particles passed straight through the gold foil. Conclusion: Conclusion: Most of the atom’s volume is empty space.

28. Observation: Observation: A few alpha particles were deflected at an angle or bounced back. A few alpha particles were deflected at an angle or bounced back. Conclusion: Conclusion: Atoms have a very small, dense positively charged nucleus. Atoms have a very small, dense positively charged nucleus.

30. Nucleus is extremely small compared to the size of the atom as a whole. Nucleus is extremely small compared to the size of the atom as a whole. Deflections happened rarely (1/8000). Deflections happened rarely (1/8000).

31. The Nuclear Model Rutherford’s Model is called the “Nuclear Model”

32. Comparison to Thomson Positively charge contained in nucleus. Positively charge contained in nucleus. Negatively particles scattered outside nucleus. Negatively particles scattered outside nucleus. Not dispursed evenly. Not dispursed evenly.

33. http://chemmovies.unl.edu/ChemAnime/R UTHERFD/RUTHERFD.html http://chemmovies.unl.edu/ChemAnime/R UTHERFD/RUTHERFD.html http://chemmovies.unl.edu/ChemAnime/R UTHERFD/RUTHERFD.html http://chemmovies.unl.edu/ChemAnime/R UTHERFD/RUTHERFD.html http://chemmovies.unl.edu/ChemAnime/R UTHERFD/RUTHERFD.html http://chemmovies.unl.edu/ChemAnime/R UTHERFD/RUTHERFD.html http://chemmovies.unl.edu/ChemAnime/R UTHERFD/RUTHERFD.html http://chemmovies.unl.edu/ChemAnime/R UTHERFD/RUTHERFD.html http://youtu.be/wzALbzTdnc8 http://youtu.be/wzALbzTdnc8 http://youtu.be/wzALbzTdnc8 http://youtu.be/XBqHkraf8iE http://youtu.be/XBqHkraf8iE http://youtu.be/XBqHkraf8iE

34. Niels Bohr (1913) Came up with the “Planetary Model” Came up with the “Planetary Model”

35. Bohr’s Theory Electrons circle nucleus in specific energy levels or “shells”. Electrons circle nucleus in specific energy levels or “shells”. The higher the “energy level” the higher the electron’s energy. The higher the “energy level” the higher the electron’s energy.

36. Energy Levels Different energy levels can contain different numbers of electrons. Different energy levels can contain different numbers of electrons.

37. How many per level? n = the number of the energy level n = the number of the energy level 2n 2 = the total number of electrons an energy level can hold. Ex: Level 3 can hold 2(3) 2 = 18 electrons

38. Draw a Bohr Atom Ex: The Fluorine Atom (F) Ex: The Fluorine Atom (F) Protons = 9 Protons = 9 Neutrons = 10 Neutrons = 10 Electrons = 9 Electrons = 9 How many energy levels do you draw? How many energy levels do you draw? How many electrons in each level? How many electrons in each level?

39. Draw a Bohr Ion They only difference is that one or more electrons gets added or taken out of the outer energy level. They only difference is that one or more electrons gets added or taken out of the outer energy level. Ex: The Magnesium Ion (Mg +2 ) Ex: The Magnesium Ion (Mg +2 ) Protons = 12 Protons = 12 Neutrons = 12 Neutrons = 12 Electrons = 10 Electrons = 10

40. (+) Ions (cations) (+) ions are smaller Lost electron(s)

41. (-) Ions (anions) (-) ions are larger Gained electron(s)

42. How Did Bohr Come Up With His Model? Studied the spectral lines emitted by various elements (especially Hydrogen) Studied the spectral lines emitted by various elements (especially Hydrogen)

43. What are Spectral Lines? Energy gets absorbed by an atom causing it to emit a unique set of colored lines. Energy gets absorbed by an atom causing it to emit a unique set of colored lines. Used to identify what elements are present in a sample. (elemental “Fingerprint”) Used to identify what elements are present in a sample. (elemental “Fingerprint”)

44. Spectral Lines are Different for Each Element

45. http://www.mhhe.com/physsci/chemistry/ essentialchemistry/flash/linesp16.swf http://www.mhhe.com/physsci/chemistry/ essentialchemistry/flash/linesp16.swf

46. What Causes Spectral Lines? Jumping Electrons!!

47. Jumping Electrons Electrons normally exist in the lowest energy level possible called the “ground state”. (stable) Electrons normally exist in the lowest energy level possible called the “ground state”. (stable) “Ground state” e - configurations are written on the periodic table for each element. “Ground state” e - configurations are written on the periodic table for each element. Ex: Aluminum is 2-8-3 Ex: Aluminum is 2-8-3 Calcium is 2-8-8-2

48. An Electron Gets “Excited” Electrons can absorb a photon of energy and “jump up” to a higher energy level farther from the nucleus. This is called the “excited state”. (unstable)

49. Jumping Electrons They quickly “fall back down” to the ground state. (stable) They quickly “fall back down” to the ground state. (stable) They emit a photon of energy that corresponds to how far they jumped. They emit a photon of energy that corresponds to how far they jumped.

50. This photon of energy is seen as a spectral line! This photon of energy is seen as a spectral line! Each spectral line corresponds to a specific photon of energy that is released. Each spectral line corresponds to a specific photon of energy that is released.

51. REMEMBER Absorb Energy Jump Up Emit Energy Fall Down

53. Electromagnetic Spectrum Spectral lines can come from all areas of the EM Spectrum. Spectral lines can come from all areas of the EM Spectrum. Lines of visible colors make up only a small part of the spectrum. Lines of visible colors make up only a small part of the spectrum.

54. EM waves carry different amounts of energy based upon their wavelength and frequency. EM waves carry different amounts of energy based upon their wavelength and frequency. Which wave has higher energy?

56. http://www.upscale.utoronto.ca/PVB/Harris on/BohrModel/Flash/BohrModel.html http://www.upscale.utoronto.ca/PVB/Harris on/BohrModel/Flash/BohrModel.html http://www.upscale.utoronto.ca/PVB/Harris on/BohrModel/Flash/BohrModel.html http://www.upscale.utoronto.ca/PVB/Harris on/BohrModel/Flash/BohrModel.html

57. Calculating the Energy of a Spectral Line STEP 1: If you know the wavelength of the spectral line you can find it’s frequency. c = λ x ү c = the speed of light = 3 x 10 8 meters/sec λ = wavelength (in meters) ү = frequency of the wave

58. Calculating the Energy of a Spectral Line STEP 2: Using the frequency find the energy of the line (in Joules) E = h x E = h x ү E = energy in Joules h = Planck's constant = 6.63 × 10 -34 kg x m 2 / sec ү = frequency of the wave

59. Electron Cloud Model

60. Sometimes called: Sometimes called: Wave Mechanical Model Quantum Mechanical Model Orbital Model Charge Cloud Model

61. How is it Different from the Planetary Model? Heisenberg’s Uncertainty Principle: It is impossible to know the exact location and momentum of an electron at the same time. We can’t tell exactly where an electron is!! We can’t tell exactly where an electron is!!

62. Electrons exist in “orbital clouds” Electrons exist in “orbital clouds” The denser the region of the cloud the higher the probability of finding an electron there. The denser the region of the cloud the higher the probability of finding an electron there.

63. http://youtu.be/45KGS1Ro-sc http://youtu.be/45KGS1Ro-sc http://youtu.be/45KGS1Ro-sc

64. How are Electrons Organized? Electron Hotel

65. Energy Levels (1-7) Electrons can be at different distances from the nucleus. Electrons can be at different distances from the nucleus. Energy Levels Energy Levels 1 23 4 5 6 7 Lowest energyHighest energy Closest to nucleusFarthest from Nucleus

66. Sublevels (s, p, d, f) Each energy level can have a certain number of sublevels. Each energy level can have a certain number of sublevels. Energy LevelSublevels Possible 1s 2s, p 3s, p, d 4s, p, d, f 5s, p, d, f, (g) 6s, p, d, f, (g, h) 7s, p, d, f, (g, h, i)

67. Energy of Sublevels Sublevels have different levels of energy. Sublevels have different levels of energy. spdf Lowest energyHighest energy

68. Orbitals in Sublevels Each sublevel contains a different number of orbitals. Each sublevel contains a different number of orbitals. A maximum of two electrons can exist in an orbital. A maximum of two electrons can exist in an orbital. Sublevel# of OrbitalsMax e - in Sublevel s12 e - s12 e - p36 e - d510 e - f714 e -

69. Electron Spin Pauli Exclusion Principle: Pauli Exclusion Principle: In order for two electrons to occupy the same orbital, they must have opposite spins. Electrons in an orbital spin in opposite directions Electrons in an orbital spin in opposite directions

70. Shapes of Orbitals Orbitals come in different shapes and sizes. Orbitals come in different shapes and sizes. They are the region of highest probability of finding an electron. They are the region of highest probability of finding an electron.

71. s Orbital Probability cloud has a spherical shape

72. p Orbitals (p x, p y, p z ) “Dumbell” shape Three p orbitals can exist, on the x, y, z axis in space

73. d Orbitals Five possible d orbitals exist Five possible d orbitals exist

74. f Orbitals Seven possible f orbitals exist Seven possible f orbitals exist