GE 150 Astronomy

1. GE 150 Astronomy Week #7 February 26,28 2013

2. The Electromagnetic Spectrum

3. Light Astronomers’ primary tool in learning about the Universe is electromagnetic radiation (i.e. all forms of light)

4. Isaac Newton showed: white light could be split into component colors with a prism and then recombined into white light with a lens White Light is a composite of all colors White Light

5. Basic Properties of Light • Light is radiant energy. • Speed – c = 300,000 km/sec • Has characteristics of both a wave and a particle

6. The colors are determined by the wavelength of light - Denoted by the Greek letter “lambda” . Measured in: Micron (μm) – 10-6m Nanometer (nm) – 10-9m Angstrom (Å) – 10-10m Basic Properties: Wavelength

7. Wavelengths of White Light long wavelengths RO Y G B I V short wavelengths Visible spectrum

8. Frequency and wavelength are related by: Basic Properties: Frequency • Frequency - how fast successive crests pass by a given point - denoted by the Greek letter “nu”  • Measured in • Hertz (Hz) = 1 cycle/sec ‘c’ is the speed of light.

9. his Planck’s constant Basic Properties: Energy Carried by Light Different colors of light have different amounts of energy

10. Light • Early telescopic observations relied on visible spectrum (i.e. the colors that the human eye can see) but today Astronomy uses the full electromagnetic spectrum

11. The Electromagnetic Spectrum

12. Consider a Class of Astronomy students as Seen in Different Wavelengths of Light!

13. Consider Orion as Seen in Different Wavelengths of Light!

14. Observations at other wavelengths are revealing previously invisible sights Infrared UV Map of Orion region Ordinary visible

15. But, what is light? Wave or Particle? • In the 17thCentury, Isaac Newton argued that light was composed of little particles while Christian Huygens suggested that light travels in the form of waves. • In the 19th and 20th Century Maxwell, Young, Einstein and others were able to show that Light behaves both like a particle and a wave depending on how you observe it.

16. Wave Interference – the 2 extremes Two waves “in phase” with each other Two wave completely “out of phase” with each other

17. Thomas Young’s interference experiment

18. Thomas Young’s interference experiment

19. Scottish physicist James Clerk Maxwell showed mathematically in the 1860s that light must be a combination of electric and magnetic fields.

20. The Nature of Light • As a particle (photons) … • Quantized - photons have very specific energies, related to its frequency • Wave theory (incorrectly) says its energy is a measure of intensity (brightness for light) • Photons strike the light sensors (or your eye) like a very small bullet - how your phone’s camera works

21. Photoelectric effect: on certain metals, shining light causes the atoms to eject electrons In 1905 Einstein explained it using light as particles (photon) and correctly predicted the energy of the ejected electrons based on the photon energy Ephoton= hnorEphoton = hc/l photon e-

22. But, where does light actually come from? Light comes from the acceleration of charged particles (such as electrons and protons)

23. But, where does light actually come from? electron Accelerating charges produce light – electromagnetic radiation!

24. Interaction of Light, Atoms and Energy

25. Atoms: Historical Perspective • Atoms first proposed by Greeks around 500 BCE • Not our concept of an atom – no protons, neutrons, electrons • More philosophical than scientific • These “atoms” couldn’t be broken down any further

26. Atoms: Historical Perspective • JJ Thompson • “plum pudding” model – 19th Century • Arose after the discovery of the electron • Electrons floating around in a positively charged soup

27. Atoms: Historical Perspective • Ernest Rutherford • in early 1909 bombarded gold foil with helium nuclei (aka alpha particles) and watched them ricochet • demonstrated atoms had a dense, charged core • Also helped develop the orbital theory of the atom • His model didn’t explain electron-structure

28. Classic “solar system” model of an atom: a small, dense nucleus (containing protons and neutrons) surrounded by electrons- Proposed by Niels Bohr 1913

29. The size of the nucleus is about 10-15m. The first electron orbits out at 10-10 m from the center of the atom –size of the electron orbit is 100,000 times greater than the size of the nucleus Atoms are mostly empty space

30. The electron should be thought of as a distribution or cloud of probability around the nucleus that on-average behave like a point particle on a fixed circular path Today’s view of the Atom

31. Atoms: Historical Perspective • Experiments in early 20th century found electrons can only orbit atoms at precisely defined levels • Discrete or Quantized energies • the beginnings of Quantum Physics

32. Review of the Atom Electron orbitals are ‘quantized’ – that is, they exist only at very specific energies • The lowest energy orbital is called the ground state

33. Quantized Electron Orbitals electronic transitions - move an electron from one orbital to another

34. Schematic of Hydrogen Atom (1 proton, 1 electron) Note the spacing between levels and how they decrease!

35. Photons (light-waves) are emitted from an atom when an electron moves from a higher energy level to a lower energy level Nucleus

36. Photons (light-waves) can also be absorbed by an atom when an electron moves from a lower energy level to a higher energy level Nucleus

37. What determines the energy Levels of the orbitals • The number of protons (atomic number) in a nucleus determines what element a substance is. • Also the same number of electrons as protons • Since each element has a unique number of protons, electron energy levels are also unique

38. What determines the energy Levels of the orbitals The energy (light) emitted or absorbed by an atom during an electronic transition indicates what element the atom belongs to, even from millions of light years away! Example: Absorption and emission of a photon in a Hydrogen Atom

39. Absorption An electron absorbs a photon of the exact energy needed to raise it an orbital

40. Emission • If an electron drops from one orbital to a lower one, it must first emit a photon with the same amount of energy as the orbital energy difference. E = h

41. Atomic Spectra: Identifying Atoms by Their Light

42. Spectroscopy The study of light emitted or absorbed by an object at various wavelengths to determine its composition and physical state (e.g. temperature)

43. How a Spectroscope Works • A narrow slit focuses the light • Agrating or a prism splits the light into its component colors

44. Each chemical element produces its own unique set of spectral lines when it is excited

45. Emission spectrum of hydrogen

46. Emission spectrum of hydrogen This spectrum is unique to hydrogen Like a barcode The red line is called ‘H alpha’ written: H

47. Which part of the spectrum has lower energy than visible? Which part of the spectrum has higher energy than visible?

48. Different atom, different spectrum Every element has its own spectrum. Note the differences between hydrogen and helium spectra below.

49. Absorption Spectra What if, instead of hot hydrogen gas, we had a cloud of cool hydrogen gas between us and a star?

50. Summary - Types of Spectra • continuous spectrum The source emits light that is continuous and all colors are present • emission-line spectrum A hot, thin gas will emit characteristic frequencies of light. • absorption line spectrum A cool gas will absorb light behind it at a characteristic frequency