1. SCH4U http://www.youtube.com/ watch?v=-d23GS56HjQ G de 12 Cmistr Ra He Y Y

2. Dalton’s Theory  Matter is made up of indestructible atoms.  Law of definite proportions:  Elements combine in a characteristic ratio  Law of multiple proportions:  Some elements have more than one combining capacity  Law of conservation of mass:  Atoms cannot be created nor destroyed

3. Thomson’s Theory  “The Raisin Bun” model:  + and – charges are mixed together  Gave us electrons  Atoms can gain or lose electrons to form ions  Said that the identity of an element was based on its number of electrons

4. Rutherford’s Model  Atoms have a tiny nucleus which contains positive & neutral charges and makes up the majority of the mass of the atom  Electrons are negative and occupy most of the volume of the atom.  Protons tell us the identity of the element

7. Atoms and Isotopes Isotopes  Have the same number of protons and electrons but have different amounts of neutrons.  Radioisotopes – give off radioactivity when they decay

8. Rutherford Model – Planetary Model of the Atom Protons Neutrons Electrons

9. Particle Mass (kg) LocationCharge Proton (p + ) 1.673 x 10 -27 Nucleus+1 Electron (e - ) 9.109 x 10 -31 Orbitals outside nucleus Neutron (n 0 ) 1.675 x 10 -27 Nucleus0

10. Representing Atoms X Z A

11. Problems - Revisited  SPIRAL DEATH!!!!

12.  To solve this problem… we need a little bit more of an insight into two phenomena:  LIGHT  ENERGY

13. Light is a Wave! Huygens, Newton

14. Light is a Particle! (The Photoelectric Effect) The ejection of electrons from a metal surface when light strikes it Certain types of light cause ejection, others don’t

15. Max Planck Spectrum of Radiated energy and intensity Quantum: unit or package of energy (plural quanta) Energy is quantize – can only have allowed values

16. Planck Equation  Energy is equal to the frequency of the radiation times Planck’s constant (h)  h = 6.64×10 -34 J∙s  Energy is QUANTIZED – it comes in packets and the smallest packet is equal to Planck’s constant  Only multiples of this number are allowed – nothing more

17. Photons  By extension, light is also a quantize, since it is a type of energy  Photon: unit of light energy  Or particles of light energy  (Used to describe the photoelectric effect)

18. Homework  Page 142 #1-7

19. Bohr’s Model of the Atom  Limitations of the Rutherford Model  Electrons orbiting around a nucleus should lose energy and spiral into the nucleus  Electrons should be attracted to proton and collapse in to the nucleus  SPIRAL DEATH

20. Atomic Spectra  Continuous Spectrum: an emission spectrum that contains all the wavelengths of light in a specific region of the electromagnetic spectrum  Line Spectrum: emission spectrum that contains only specific wavelengths characteristic of the element being studied

23. Hydrogen Emission Spectrum

24. Reason?

25. Different for Each Element

26. Bohr’s Postulates  First Postulate:  e- do not radiate energy as they orbit the nucleus. Each orbit corresponds to a state of constant energy (called stationary state).  Basically energy states (or levels)

27.  Second Postulate:  e- can change their energy only by undergoing a transition from one stationary state to another  Basically, give the e- a quantum of energy and it’ll jump up to the next energy level, when it loses the quantum it falls back down, releasing a photon

28. Bohr-Rutherford Model

29. Successes and Failures of the Bohr Model  Works well at predicting properties and periodicity of the elements  Problem: everything was a little bit off after Hydrogen.

30. Trends in the Periodic Table  Atomic radius  Ionization Energy  Electron Affinity  Electronegativity

31. Homework

32. THE QUANTUM MECHANICAL MODEL OF THE ATOM

33. And now for something completely different…

34. Quantum Mechanics  The application of quantum theory to explain the properties of matter, particularly electrons in atoms

35. Schrodinger’s Standing Waves  Louis De Broglie developed a theory that matter can have wave-like properties  Schrodinger extended this theory to electrons bound to a nucleus  Postulated that electrons resembled a standing wave  Certain orbitals exist at whole wavelengths of electron vibrations

37. Orbitals - Redefined  Orbital: region around the nucleus where there is a high probability of finding an electron  As per wave model of Schrodinger – because things are vibrating

38. Heisenberg Uncertainty Principle

39.  Heisenberg studied statistics and developed matrix algebra  Developed a statistical approach to explaining how electrons works and realized…  IT IS IMPOSSIBLE TO KNOW THE EXACT POSITION AND SPEED OF ELECTRON AT A GIVEN TIME  At best, we can describe the probability of finding it at a specific place

40.  Wave functions: the mathematical probability of finding an electron in a certain region of space  Wave functions give us:  Electron probability densities: the probability of finding an electron at a given location, derived from wave equations

43. Homework

44. Quantum Numbers  Quantum Numbers: numbers that describe the quantum mechanical properties (energies) of orbitals  From the solutions to Schrodinger’s equation  The most stable energy states is called the ground state

45. Principal Quantum Number (n)  Integer number (n) used to level the main shell or energy level of the electron  Describes size and energy of the atomic orbital  Increase number = increase energy, bigger

46. Secondary Quantum Number, l  Describes the shape of the orbital within each shell  Each energy level contains several sublevels  Relates to the shape of the orbital  Can be any integer from 0 to (n-1)

47. Values of l Value01234 Letter Used spdfg Namesharpprincipaldiffusefundamental

48.  Each orbital is given a code:  Example  If n = 1, l = 0 then we call it a 1s orbital  If n = 3, l = 2 then we call it a 3d orbital

49. Magnetic Quantum Number, m l  Describes the orientation of the orbital in 3- space  Can be whole number integers from – l to + l  Example: if l = 1, then m l can be -1, 0, +1  There are 3 possible p orbitals  p x, p y, and p z

50.  What are possible values for m l if l is:  0  1  2  3

52. Spin Quantum Number  Electrons are basically little magnetics spin around when placed in magnetic fields, they can have spin ‘up’ or spin ‘down’  m s can be either +1/2 or – 1/2

53. Homework

54. Electron Configurations and Energy Level Diagrams  The four quantum numbers tell us about the energies of electrons in each atom  Unless otherwise stated were are talking about ground state energies

55. Energy Diagrams  Describe how electrons fill orbitals using quantum numbers  Electrons fill the lowest energy level orbitals first  Each shell is (for the most part) filled before moving to higher shells

58. Rules  Use circles (or boxes) to represent each orbital in any given energy level and arrows for electrons  Unoccupied circles imply that there are no electrons in it  A circle can have at most two electrons in it; only if the arrows are pointing in opposite directions

59. Rules  Pauli exclusion Principle: no two electrons can have the same 4 quantum numbers. Electrons in the same orbital can’t have the same spin  Hund’s Rule: One electron occupies each of several orbitals in the same energy level before a second can occupy the same orbital  Aufbau Principle: each electron is added to the lowest energy orbital avaible

60. Practice  H, B, C, Ne  Mg, P, Ar  Ca, Mn, Zn, Ge, Kr

61. Electron Configurations  Condensed versions of orbital diagrams and not in  Write the electron configuration for each of the atoms above

62. Exceptions to the Rules