Adv chem chapt 7

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Transcript of Adv chem chapt 7

  • 1.ATOMIC STRUCTURE AND PERIODICITY Advanced Chemistry 7

2. ELECTROMAGNETIC RADIATION 3. ELECTROMAGNETIC RADIATION

  • One of the ways that energy travels through space is by electromagnetic radiation.
  • light from the sun
  • X-rays
  • microwaves

4. WAVE CHARACTERISTICS

  • Wavelength () is the distance between two consecutive peaks or troughs in a wave.
  • Frequency () is the number of waves (cycles) per second that pass a given point in space.
    • units hertz or waves/sec (s -1 )
  • Speed (c) all types of electromagnetic radiation travel at the speed of light.
    • 2.9979 x 10 8m/s
    • c =

5. ELECTROMAGNETIC RADIATION 6. ELECTROMAGNETIC RADIATION 7. THE NATURE OF MATTER 8. WAVE AND PARTICLE DUALITY

  • Planck found that matter could only absorb or emit energy in whole number multiples of the quantityh .
  • his Plancks constant = 6.626 x10 -34J s
  • E = h
  • Transfer of energy is not continuous but is quantized and can occur only in discrete amounts calledquantum .Thus energy has particle properties as well as wave properties.

9. EINSTEIN 10. WAVE AND PARTICLE DUALITY

  • Einstein proposed that electromagnetic radiation was also quantized and could be viewed as a stream of particles calledphotons .
  • E photon=hv=hc/

11. THE PHOTOELECTRIC EFFECT

  • The photoelectric effect refers to the phenomenon in which electrons are emitted from the surface of a metal when light strikes it.
  • No electrons are emitted by a metal below a specific threshold frequency ( v o )
  • For light with frequency lower than the threshold frequency, no electrons are emitted regardless of intensity of the light.

12. THE PHOTOELECTRIC EFFECT

  • For light with frequency greater than the threshold frequency, the number of electrons emitted increases with the intensity of the light.
  • For light with frequency greater than the threshold frequency, the kinetic energy of the emitted electrons increases directly with frequency of the light.

13. THE PHOTOELECTRIC EFFECT

  • These observations can be explained by assuming that electromagnetic radiation is quantized (consists of photons), and that the threshold frequency represents the minimum energy required to remove the electron from the metals surface.
  • Minimum energy required to remove an electron = E o=hv o
  • KE electron= mv 2=hv hv o

14. PLANCK AND EINSTEIN CONCLUSIONS

  • Energy is quantized.It can occur only in discrete units called quanta.
  • Electromagnetic radiation, which was previously thought to exhibit only wave properties, seems to show certain characteristics of particulate matter as well.This phenomenon is sometimes referred to as the dual nature of light.

15. WAVE PARTICLE DUALITY

  • The main significance of the equation E = mc 2is that energy has mass.
  • m = E/c 2

16. LOUIS DE BROGLIE (1892-1987)

  • Since light which previously was thought to be purely wavelike, was found to have certain characteristics of particulate matter.But is the opposite also true?Does matter have that is normally assumed to be particulate exhibit wave properties?

17. LOUIS DE BROGLIE (1892-1987)

  • de Broglies equation allows us to calculate the wavelength for a particle:

18. DE BROGLIES PROOF 19. LOUIS DE BROGLIE (1892-1987)

  • Conclusion:Energy is really a form of matter, and all matter shows the same types of properties.All matter exhibits both particulate and wave properties.

20. THE ATOMIC SPECTRUM OF HYDROGEN 21. SPECTRUM

  • Acontinuous spectrumresults when white light passes through a prism and all wavelengths (colors) are shown.
  • Anemission spectrumproduces only a few lines of color that is limited to discrete wavelengths produced by an atom.This is called aline spectrumand is specific to each atom.

22. HYDROGEN LINE SPECTRUM

  • The significance of the line spectrum is that it indicates that only certain energies are allowed for the electron in the hydrogen atom.In other words the energy of the electron in the hydrogen atom is quantized

23. HYDROGEN LINE SPECTRUM 24. THE BOHR MODEL 25. NIELS BOHR

  • Bohr developed a quantum model for the hydrogen atom that allowed for only specific energy levels around the atom that corresponded with specific radii.

26. NIELS BOHR (1885-1962)

  • The most important equation to come from Bohrs model is the expression for the energy levels available to the electron in the hydrogen atom.
  • Z is the nuclear charge, n is the energy level.

27. NIELS BOHR (1885-1962)

  • The most important equation to come from Bohrs model is the expression for the energy levels available to the electron in the hydrogen atom.
  • the negative sign calculates a lower energy closer to the atom, not the radiation of negative energy.

28. EXAMPLE

  • What is the change in energy if an electron in level 6 (excited state) returns to level 1 (ground state) in a hydrogen atom?
  • n i =6; n f =1; Z=1 (hydrogen nucleus contains a single proton)

29. EXAMPLE

  • What is the change in energy if an electron in level 6 (excited state) returns to level 1 (ground state) in a hydrogen atom?

30. EXAMPLE

  • E=E f E i= E 1 E 6 =-2.117 x 10 -18 J
  • The negative sign for the change in energy indicates that the atom has lost energy and is now more stable. This loss of energy produces a photon.

31. EXAMPLE

  • What is the corresponding wavelength for the energy produced from the electron jump?
  • E = -2.117 x 10 -18 J

9.383x10 -8m 32. BOHR MODEL CONCLUSIONS

  • The model correctly fits the quantized energy levels of the hydrogen atom and postulates only certain allowed circular orbits for the electrons.
  • As the electron becomes more tightly bound, its energy becomes more negative relative to the zero-energy reference state.As the electron is brought closer to the nucleus, energy is released from the system.

33. BOHR MODEL CONCLUSIONS 34. BOHR MODEL CONCLUSIONS

  • The energy levels calculated by Bohr closely agreed with the values obtained from the hydrogen emission spectrum but does not apply well to other atoms.The Bohrs model is fundamentally incorrect but is very important historically because it paved the way for our current theory of atomic structure.

35. THE QUANTUM MECHANICAL MODEL OF THE ATOM 36. QUANTUM MECHANICS

  • Quantum Mechanics or Wave Mechanics were developed by three physicists: Heisenberg, de Broglie, and Schrodinger.
  • Emphasis was given to the wave properties of the electron.
  • The electron bound to the nucleus behaves similar to a standing wave.

37. QUANTUM MECHANICS

  • Like a standing wave, electrons can travel in patterns that allow for a common node.In other words, wave patterns around the nucleus must be in whole number wave patterns.But their exact movement is not known.

38. HEISENBERG UNCERTAINTY PRINCIPLE

  • There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time.This limitation is small for large particles but substantial for electrons.

39. PROBABILITY DISTRIBUTION

  • A probability distribution is used to indicate the probability of finding an electron in a specific position.
  • Electron density map
  • Radial probability distribution

40. PROBABILITY DISTRIBUTION

  • For the hydrogen 1s orbital, the maximum radial probability occurs at a distance of 5.29x10 -2 nm or .529 from the nucleus.This is the exact radius of the innermost orbit calculated in the Bohr Model.
  • The definition most often used by chemists to describe the size of the hydrogen 1s orbital is the radius of the sphere that encloses 90% of the total electron probability

41. QUANTUM NUMBERS 42. QUANTUM NUMBERS

  • Each orbital is characterized by a series of numbers called quantum numbers, which describe various properties of an orbital:
  • Principal quantum number (n) - has integral values : 1,2,3,4.It describes the size and energy of the orbital.Energy Level

43. QUANTUM NUMBERS

  • Angular momentum quantum number ( l ) has integral values from 0 to n-1.This is related to shape of the atomic orbitals.Sublevel
    • l=0 is s
    • l=1 is p
    • l=2 is d
    • l=3 is f
    • l=4 is g

44. QUANTUM NUMBERS

  • Magnetic quantum number (m l )has values betweenland l, including 0.The value of m lis related to the orientation of the orbital in space.Axis designation

45. QUANTUM NUMBERS

  • Electron spin quantum number (m s )-can only have one of two values, +, -.Electrons can spin in one of two opposite directions.

46. QUANTUM NUMBERS

  • In a given atom no two electrons can