1. Atomic Physics – Part 1 To accompany Pearson Physics PowerPoint presentation by R. Schultz robert.schultz@ei.educ.ab.ca

2. 15.1 The Discovery of the Electron Cathode rays – major item of interest near the end of the 1800’s – what were they? Emr? Charged particles? Neutral particles? Path deflection caused by both electric (parallel to field) and magnetic (perpendicular to field) fields

3. 15.1 The Discovery of the Electron Conclusion: negatively charged particles J.J. Thomson, 1897, measured charge to mass ratio of a cathode particles using the apparatus shown: particles accelerated from C to A focused by B deflected by magnetic field from coil straightened by electric field between D and E

4. 15.1 The Discovery of the Electron Thomson found q/m for cathode ray particles to be 1.76 x 10 11 C/kg This is close to 2000 x q/m for a hydrogen ion Thomson’s conclusion: cathode-ray particles (today called electrons) are approximately 1/2000 the mass of a hydrogen ion (today called protons)

5. 15.1 The Discovery of the Electron Since Thomson didn’t know the charge or mass of these particles couldn’t use to determine their speed He balanced so that particle path was undeflected

6. 15.1 The Discovery of the Electron Then, Using only a magnetic field,

7. 15.1 The Discovery of the Electron Examples: Practice Problem 1, page 756 Practice Problem 1, page 758

8. 15.1 The Discovery of the Electron The Mass Spectrometer uses the principles behind Thomson’s experiment to measure the charge to mass ratio of compounds Read Then, Now and Future, page 759

9. 15.1 The Discovery of the Electron The Thomson Raisin-bun model of the atom Recalling that Thomson concluded electrons were approximately 1/2000 the mass of equivalent amount of positive charge, he concluded that atom was a mass of (+) charge, taking up almost total volume of atom, with tiny, near massless, electrons embedded in it Like raisins in a bun – Thomson Raisin Bun Model

10. 15.1 The Discovery of the Electron Thomson quotation Check and Reflect page 760, questions 3, 7 SNAP booklet, page 274, questions 1, 3, 4, 7, 9, 11, 13, 16

11. 15.2 Quantization of Charge Millikan Oil-Drop Experiment oil droplets sprayed into chamber fall into region of strong electric field voltage across plates adjusted until gravity is balanced Millikan’s actual apparatus still at CalTech

12. 15.2 Quantization of Charge Because of air resistance actual analysis is more complex than that presented here We will stick with the analysis required for this course

13. 15.2 Quantization of Charge Example: Practice Problem 1, page 763

14. 15.2 Quantization of Charge In some problems you will need to use the density of the oil and radius of droplet to find its mass mass = density x volume

15. 15.2 Quantization of Charge SNAP booklet, page 280, questions 1, 2, 3, 7, 8, 10

16. 15.3 The Discovery of the Nucleus Discuss QuickLab 15-3, page 766 The Rutherford Experiment

17. 15.3 The Discovery of the Nucleus Results most alpha particles (helium nuclei) passed through gold foil with scattering of 1º or less A small number were scattered backwards (at angles greater than 140º)

18. 15.3 The Discovery of the Nucleus Not predicted by Thomson raisin bun model Because of low density of mass and positive charge in Thomson Raisin Bun model, alpha particles ( α particles) expected to blast right through bun Even direct collisions with electrons (raisins) should cause little scattering

19. 15.3 The Discovery of the Nucleus Conclusion: Positive charge and mass not distributed throughout whole atom, but concentrated in a tiny fraction of total volume – the nucleus Atom radius approximately 10 -10 m Nucleus radius approximately 10 -14 m Volume to volume – like an ant in a football field Atom mostly empty space!

20. 15.3 The Discovery of the Nucleus Rutherford’s famous quotation – the scattering “was almost as incredible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you!”

21. 15.3 The Discovery of the Nucleus Rutherford’s Planetary Model of the atom Review example 15.5, page 769 Do and discuss Check and Reflect, page 770, questions 1, 2, 4, and 5 Note if this diagram was to scale, nucleus would be too small to be visible

22. 15.4 The Bohr Model of the Atom Problem with Rutherford model: Recall that accelerating charges produce electromagnetic waves Electrons orbiting nucleus undergo continuous centripetal acceleration Atom would give off emr, lose energy, and electrons spiral into nucleus in small fraction of a second!

23. 15.4 The Bohr Model of the Atom Fraunhofer, 1814, noticed dark lines in the solar spectrum Further study of the spectra led to the following generalizations:

24. 15.4 The Bohr Model of the Atom Hot dense material (e.g. a glowing solid) produces continuous spectrum with no bright or dark lines Hot dense material (e.g. a glowing solid) produces continuous spectrum with no bright or dark lines Hot gas at low pressure: emission, bright line spectrum, characteristic of element Hot gas at low pressure: emission, bright line spectrum, characteristic of element

25. 15.4 The Bohr Model of the Atom Cool gas at low pressure with light shone through it, absorption, dark line spectrum Cool gas at low pressure with light shone through it, absorption, dark line spectrum Dark lines of absorption spectrum occur at the same frequencies as the bright line spectrum for the same gas

26. 15.4 The Bohr Model of the Atom Fraunhofer lines on solar spectrum later realized to be absorption spectra of all of the gases in the cooler outer atmosphere of the Sun Elements identified by comparing individual elements’ spectra with lines on the solar spectrum

27. 15.4 The Bohr Model of the Atom Hydrogen’s emission spectrum had been analyzed by Balmer, a mathematician, in 1885 Balmer noticed a pattern and formulated an equation to predict wavelengths emitted you won’t do calculations with this formula on the diploma exam

28. 15.4 The Bohr Model of the Atom Bohr realized that emitted wavelengths of light were due to differences between quantized energy levels in hydrogen atom He postulated mathematically that: Electrons were allowed to orbit nucleus at certain allowed radii (no others), called stationary states, without emission of radiation Electrons were allowed to orbit nucleus at certain allowed radii (no others), called stationary states, without emission of radiation

29. 15.4 The Bohr Model of the Atom Electrons in each stationary state orbit had a fixed amount of energy – the closer to the nucleus, the lower the energy (energy was quantized) Electrons in each stationary state orbit had a fixed amount of energy – the closer to the nucleus, the lower the energy (energy was quantized) Electrons can move from one stationary state to another by giving off or absorbing energy in the form of EMR (remember the spectrum of hydrogen) Electrons can move from one stationary state to another by giving off or absorbing energy in the form of EMR (remember the spectrum of hydrogen)

30. 15.4 The Bohr Model of the Atom Allowed radii for hydrogen: Allowed energy for hydrogen: Don’t need to know for test -13.6 eV is ground state (n=1) for hydrogen n=2 and higher called “excited states”

31. 15.4 The Bohr Model of the Atom

32. View emission spectrum of hydrogen hydrogen tube metre stick diffraction grating view hydrogen tube through grating see spectral lines on top of metre stick

33. 15.4 The Bohr Model of the Atom Discuss Check and Reflect, page 780, questions 2, 4, 8, 10 You should know in a non-quantitative manner, the Bohr model of the atom: what are the basic features, how does it get around the difficulty in the Rutherford model, how does it explain the spectrum of hydrogen You should be familiar with the three types of spectra, and how to produce each

34. 15.5 The Quantum Model of the Atom Bohr’s model accurately predicts hydrogen’s spectral lines, size of the unexcited atom, and ionization energy Shortcomings: Doesn’t explain why energy is quantized or how electrons can orbit a specific radii and not emit emr Doesn’t work completely for anything but hydrogen

35. 15.5 The Quantum Model of the Atom Doesn’t explain why spectral lines are split by a magnetic field (Zeeman Effect)

36. 15.5 The Quantum Model of the Atom Recall De Broglie – previous chapter – wavelength of particles: Extend the idea – the allowed orbits for hydrogen are the right size so that each is a whole number of “standing electron waves” other orbital radii can’t exist because they won’t allow a whole # of standing waves

37. 15.5 The Quantum Model of the Atom Schrödinger, 1926, wrote a wave function (called the psi “Ψ” function) to describe electron waves Ψ 2 for a point is proportional to the probability of finding an electron at that point Orbitals – probability distributions for electrons of different energies

38. 15.5 The Quantum Model of the Atom Quantum indeterminancy: in the quantum mechanical model, electrons don’t orbit the nucleus at defined orbits, they don’t have a precise location Einstein and also Schrödinger himself, had difficulty accepting this interpretation of the model Interesting reading – The Schrödinger’s cat paradox

39. 15.5 The Quantum Model of the Atom