Presentation Transcript

1. Solid Sphere Model orBilliard Ball Model John Dalton Plum Pudding Model J.J. Thomson Planetary Model or Nuclear Model Ernest Rutherford http://mhsweb.ci.manchester.ct.us/Library/webquests/atomicmodels.htm

2. Rutherford’s model did not account for chemical properties or actual arrangement of the electrons around the nucleus • The way subatomic particles arrange themselves (mostly e-) affects properties of atoms like chemical behavior • To account for chemical behavior must look at electron configuration • Idea that came from studying behavior of light and atoms and electromagnetic radiation

3. Flaws with Rutherford’s Model, ctd. If an atom is mostly empty space, and if the electron is floating around in that empty space, then the atom should collapse (since the electron would be attracted to the nucleus) Niels Bohr proposes that e- are in circular paths around the nucleus (orbits/energy levels) Bohr Model or Orbit Model Neils Bohr

4. First we will look at properties of light and waves • Then we will look at the modifications to Rutherford’s model of the atom • Mainly Bohr and Quantum Mechanics

5. Electromagnetic Radiation • A way that energy travels through space • Ex. Sun, microwaves, x-rays, etc. • Although the forms of E are different, they have important similarities • Wavelike behavior • Travel at speed of light in vacuum c= 3x108 m/s (which is 186,000 miles/second)

6. Radiation moves like a wave  Waves characterized by two properties: wavelength () and frequency () : distance between waves (crest to crest) measured in length units: usually nm (How does this relate to m?) : number of waves passing through a fixed point in one second units: per second (1/s or s-1 or Hz)

7. How do  and  relate? Suppose this is one second-what is happening?

8. Waves move at the speed of light, c • Recall: wavelength and frequency are inversely proportional c =  = length/time (nm/s or m/s)

9. What is the wavelength, in nm, of light with a frequency of 3.91x1014 s-1? • What is the frequency of light with a wavelength of 456 nm? • Do your answers make sense?

10. Continuum of waves: shows the different forms of radiation and how they relate energetically

11. Comparing Energy Compare the energy of radio waves (~100 m) to that of UV-rays (10-4 m) What are these wavelengths in nm? What are the frequencies of these wavelengths?

12. What about Energy, E? • Proposed by Max Planck in 1900 • Energy is quantized • Comes in small, discrete, whole # packets • E is proportional to frequency and inversely proportional to wavelength: E = nh = nhc/ h=Planck’s constant = 6.626x10-34 J·s

13. Quantized Energy • Means energy comes in small whole # packets of nh is the small whole # packet of energy (n is a multiple integer) • Can only travel in increments of say 1 mph, 2mph, 3mph (not 1.5 mph) • More relevant for small objects like e- • Less noticeable as object get bigger (like car or a human)

14. In 1905, Einstein used Planck’s theory to explain the photoelectric effect. • If you shine light on a clean metal surface, electrons can be emitted (in the form of color and light) • Einstein assumed that the radiant energy hitting metal surface behaves like a stream of tiny energy packets (quanta) • Einstein calls these streams of energy photons Ephoton = h Meaning, radiant energy is quantized

15. Photoelectric Effect • To eject an e-, the light hitting the sample must have a minimum energy (threshold energy or work function). • If not at threshold energy, no light will be emitted from sample. • no matter the intensity or amount of light • This is the photoelectric effect Think about cathode ray, will it work with regular light bulb as energy source (vs. cattle prod)?

16. What is the energy of a photon whose  = 589 nm? • What is the energy per mol of the photon? Need NA

17. Dual Nature of Light • Einstein believed that light was a stream of photons (particles) • Means there is a dual nature of light • Has properties of both waves (move like waves) and particles (has mass and energy is quantized)

18. Flaws with Rutherford’s Model, ctd. If an atom is mostly empty space, and if the electron is floating around in that empty space, then the atom should collapse (since the electron would be attracted to the nucleus) In 1913, Bohr proposes that e- are in circular paths around the nucleus (orbits/energy levels) Bohr Model or Orbit Model Neils Bohr

19. Bohr’s Model • e- can only be certain distances from the nucleus • Only orbits of certain radii, corresponding to specific E, can be occupied • Distance is based on the amount of energy in the e- • e- close to the nucleus have low energy • those far from the nucleus have higher energy

20. Bohr’s Model • Energy difference between two levels = a quantum of energy

21. Bohr Model/Ladder Analogy • If you are on a low rung, do you have low or high PE? • Can your foot move from rung A to rung B? • requires energy input • What happens to PE as you move up the ladder? • Can your foot be in between two rungs? • Energy is quantized

22. Just as your foot can move between two rungs of a ladder, e- can move (jump) between allowed orbits • In the transition, energy is absorbed or released Bohr Illustration http://www.unm.edu/~astro1/101lab/lab5/lab5_C.html

23. In the transition, energy is absorbed or released • The energy is emitted or absorbed as a photon whose E = h

24. The energy of each level can be calculated using the following equation E = -(2.18x10-18 J)(1/n2) where n = energy level Ground state (GS) has n = 1 Next allowed level (ES) has n = 2 etc.

25. What is energy if n =1? • n = 2? • Notice, as n increases, E gets less negative (less favorable) • As you get further from nucleus, electron is less stable

26. Bohr said e- can jump between allowed energy levels • The energy to do this can be calculated: ∆E = -(2.18x10-18 J)(1/nf2 – 1/ni2) where ni is the energy level you start from and nf is the energy level you end up in

27. Calculate the energy needed to move an electron from GS to third ES. • Notice sign for ∆E. Does it make sense? • Calculate the energy and  to move an electron from n=5 to n=2. • What should sign of energy be? • Notice  is in visible part of spectrum. Explanation for line spectra (next slide)

29. Lines are consistent with the idea that quantization limits the possible energies allowed wikipedia

30. Bohr’s model/math could explain the behavior of atoms with only 1 e- • He could not explain the energy/colors given off by anything with more than 1e- • 1st Flaw: orbits are not circular (we learn this from quantum mechanics) • If circular, e- should lose energy as it moves and would eventually crash into + nucleus • 2nd Flaw: described the path of the electron like a large moving object. • Discredited the fact that it actually moves like a wave. (deBroglie)

31. What We Keep from Bohr • e- exist only in discrete energy levels described by quantum numbers • Energy is needed to promote e- to higher energy levels • Energy is released when e- move from high to low energy levels

32. Louis de Broglie (1924) • Electrons (particles) move like waves • goes along with Einstein’s dual nature of light • confined waves can only have certain frequencies • Only a certain # waves can pass through that space in a second • de Broglie suggested that electrons are like waves confined to a space around nucleus (orbitals) • If wavelike, they can only have certain frequencies (or energies)

33. de Broglie • The  of an e- (or any particle) depends on its mass and velocity  = h/mv mv is momentum • What is the  of an e- moving with a speed of 5.97x106 m/s? m of e- = 9.11x10-31 kg

34. If e- behaves like a wave can we determine its position? • a wave extends into space and its location is not precisely defined • Heisenberg’s Uncertainty Principle • We can’t know both the location and momentum of an e- (because of dual nature) • There will be uncertainty in position, ∆x and momentum, ∆mv ∆x · ∆mv  h/4

35. Because of dual nature and uncertainty principle… • New atomic model is based on the idea that we can calculate the energy of an electron but can only describe location in terms of probability (2) (Schrodinger) • Enter quantum mechanics

36. If e- move like waves, they have certain frequencies • These frequencies could correspond to the energy levels/orbitals

37. Orbitals • Regions around the nucleus that corresponds to specific energy levels • Electron clouds • Regions where electrons are likely to be found • Cloud is more dense in regions of higher probability • More likely near nucleus or far from nucleus? Why?

39. Solid Sphere Model orBilliard Ball Model John Dalton Planetary Model orNuclear ModelErnest Rutherford Plum Pudding Model J.J. Thomson Bohr Model or Orbit Model Neils Bohr Electron Cloud Model orQuantum Mechanical ModelLouis de Broglie & Erwin Schrodinger http://mhsweb.ci.manchester.ct.us/Library/webquests/atomicmodels.htm

40. Bohr • Circular orbits • E quantized • Treats e- as large moving object • Only works for 1e- systems Quantum Mechanics • Mathematical description of atom • Orbitals (probable location) • E quantized • Treats e- like wave