Understanding Atomic Models and Quantum Mechanics Concepts

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

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 Electrons Protons Neutrons

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 Uncertainty Principle • 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

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, ml • Describes the orientation of the orbital in 3-space • Can be whole number integers from – l to + l • Example: if l = 1, then ml can be -1, 0, +1 • There are 3 possible p orbitals • px, py, and pz

50. What are possible values for ml if l is: • 0 • 1 • 2 • 3