Chapter 6 Modern Atomic Theory

1. Chapter 6Modern Atomic Theory

2. Review… • Dalton • Thomson • Rutherford • Model doesn’t explain how the negative electron can stay in orbit and not be attracted to the positive proton

3. Electromagnetic Radiation • Light travels in • Light is a form of • Form of energy that exhibits

4. Electromagnetic Radiation • All waves can be described in 3 ways: • Amplitude – • Wavelength (l): • Frequency (n):

5. Electromagnetic Radiation • Speed of light in air: Electromagnetic radiation moves through a vacuum at speed of • Since light moves at constant speed there is a relationship between wavelength and frequency: Wavelength and frequency are inversely proportional

6. Electromagnetic Spectrum

7. Photoelectric Effect • The emission of • Albert Einstein (1905) used Planck’s equation to explain this phenomenon; • proposed that light consists of • Photon =

8. Photoelectric Effect • He (Einstein) explained that the photoelectric effect would not occur if the frequency and therefore • Analogy: • 70 cents placed in soda machine: no soda • 30 cents more and you will get your soda

9. Niels Henrik David Bohr • 1885-1962 • Physicist • Worked with Rutherford • 1912 • Studying line spectra • of hydrogen

10. Niels Henrik David Bohr • 1913 – proposed new atomic structure • Electrons exist in • Electrons

11. The Bohr Atom • Nucleus with • Electrons move in • When an electron moves from one state to another the energy lost or gained is in • Each line in a spectrum is produced when an electron moves from

12. The Bohr Atom • Model didn’t seem to work with atoms with more than one electron • Did not explain chemical behavior of the atoms

13. Now… • Light can be described as • What does this mean for the atom???

14. LineSpectrum • Elements in gaseous states give off colored light • High temperature or high voltage • Always the same • Each element is unique • Spectra

15. Line Spectrum • Ground state • Excited state

16. Line Spectrum • Electron • Color of light emitted depends on

17. Line Spectrum • Each band of color is produced by light of a different • Each particular wavelength has a definite • Each line must therefore be produced by emission of photons with

18. Line Spectrum • Whenever an excited electron • The energy of this photon is equal to the difference

19. Wave Matters… • Louis de Broglie (1924) • Proposed that electrons might have a • Used observations of normal wave activity

20. Problems… • Wave theory does not explain • Heated iron gives off heat • 1stred glow yellow glow white glow • How elements such as barium and strontium give rise to green and red colors when heated

21. Beginnings… • Max Planck (1858-1947) • Proposed that there is a fundamental restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy • Energy is released in

22. Beginnings • A quantum is a finite quantity of energy that can be gained or lost by an atom • This constant, h, is the same for all electromagnetic radiation

23. Bohr’s Equation • En = (-RH)(1/n2) • Where RH = 2.18 x 10-18J • And n = principal quantum number, 1 to infinity

24. Jumping electrons… • If an electron moves from one energy level to another, the change in energy can be determined by the following equation: • E = Ef – Ei = hν • Or simply: E = hv • Where h=6.626 x 10-34 J s

25. Then… by substitution… ( E RH ( 1 1 - = ν = ni2 nf2 h h

26. Finally… Matter waves • All moving particles • Some is apparent, some not. • De Broglie’s equation h λ = mν

27. Smart guy… • Erwin Schrodinger (1926) • Used mathematical understanding of wave behavior – devised an equation that treated electrons moving around nuclei as waves • Quantum Theory

28. Uncertainty principle • Heisenberg:

29. Quantum Theory • Describes mathematically the wave properties of electrons and other very small particles • Applies to all elements (not just H)

30. Quantum Numbers • Numbers that specify the • Principle Quantum Number: • Symbolized by n,

31. Energy Levels of Electrons • Principle energy levels • Designated by letter n • Corresponds to the • Each level divided into sublevels • 1st energy level has • 2nd energy level has • Etc.

32. Orbitals • Electrons don’t • Orbital: region in space where • Each orbital sublevel can hold

33. Orbitals Each sublevel (orbital) has a specific shape http://daugerresearch.com/orbitals/

34. Quantum Numbers • Orbital Quantum Number: • Indicates the shape of an orbital • (subshell or sublevels) • s, p, d, f Principal Quantum # Orbital Quantum # 1 2 3 4

35. Quantum Numbers • Magnetic Quantum Number: • Indicates the • Orbital position with respect to

36. Orbitron • For a full view of the different orbital shapes, visit • http://www.shef.ac.uk/chemistry/orbitron/index.html

37. Orbitals • Pauli exclusion principle: • Electrons can only spin • Shown with

38. Rules for Orbital Filling • Pauli’s Exclusion Rule • No two electrons have • Hund’s Rule • Electrons will remain 1s 2s 2p 3s 3p

39. Rules for Orbital Filling • Diagonal Rule • The order of filling once the d & f sublevels are being filled • Due to energy levels

40. Rules for Orbital Filling

41. Application of Quantum Numbers • Several ways of writing the address or location of an electron • Lowest energy levels are filled first • Electron Configuration: 12C: 32S:

42. Application of Quantum Numbers • Orbital filling electron diagram: using Hund’s rule and the diagonal rule write out the location of all electrons • See examples on whiteboard