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How atoms produce light

How atoms produce light

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How atoms produce light

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  1. How atoms produce light

  2. What is Light? • LIGHT is a form of energy • Light can be considered as a bunch of individual light “packets” called PHOTONS • Each packet has its own set of properties (wavelength, etc) • A bunch of these packets traveling together is like a ray of light.

  3. Spectroscope • A simple spectroscope has a flat prism that separates light so we can see the individual colors.

  4. “White Light” • Aim the vertical slit toward the incandescent light source. • You will see the light through the slit. • Without moving the spectroscope, drift your eyes to the right until you see the numbers on the scale. • What do you see over the numbers?

  5. Continuous Spectrum= all colorsThere are no “blank spots” in the spectrum!

  6. Why continuous spectrum? • A solid is heated…all of its atoms/molecules and their parts move really fast • Energy is given off as the atoms constantly vibrate. • Photons of all colors can be emitted. • All colors blend into “white light”

  7. Another type of spectrum • Aim the vertical slit at the overhead lights in the room. • How does this look different from the incandescent light?

  8. What kind of spectrum? • Look at the overhead lights again. Did anything change? • Are photons of ALL colors being given off by these lights? • This is called a brightline spectrum!

  9. Brightline Spectrum When only certain photons are observed, it means that only light packets of a particular type are being emitted! • Each photon has a specific energy value. • So only certain energy exchanges are happening within the heated substance. • So there must only be certain ways of changing the energy in the substance!

  10. How? • This can be explained by the movement of electrons! • We know from middle school that atoms have “layers” of electrons called energy levels. • Each energy level has electrons with a certain amount of energy in them that matches the level. • When the electrons change levels, they have to gain or lose energy to do so. • Each time they lose energy, they emit a bundle of energy. • We see that bundle as a photon!

  11. Observing Elements • You will use the spectroscopes to see what photons are given off by the elements used in the flame test. • There will be some “background” light in your spectra. • Only focus on the bright lines!

  12. Atomic Spectra/Flame Tests • When we heat up an element, its electrons gain and lose energy while it is being heated. • The photons released with each energy loss travel together as rays of light that blend into a specific color. • This light can be passed through a prism so you can see the individual colors of each type of photon.

  13. Sodium Sodium Emission Lines 400 500 600 700 Wavelength in nanometers

  14. Analyze the spectrum • How many different types of photons are in the visible range for a sodium atom emission? • Two! Both are yellow, but have different wavelengths, so are different yellows.

  15. Copper Copper Emission Lines 400 500 600 700 Wavelength in nanometers

  16. Analyze the spectrum • How many different types of photons are in the visible range for a copper atom emission? • Eight! A purple one, a blue one, and several wavelengths of greens and yellows.

  17. The individual photon colors emitted by the electrons in any atom form the “atomic emission spectrum” Also called “brightline spectrum”

  18. Conclusion • Atoms only emit photons of specific energies • WHY??

  19. All about… light

  20. LIGHT • A form of energy! • Travels in waves • Wave properties are all related • All light is part of the electromagnetic spectrum (like energy from the sun)

  21. Wave properties • Speed • Wavelength • Frequency • Energy

  22. Speed • Light travels at the speed of light (duh!) • The speed of light = 2.998 x 108 m/s • The symbol “c” stands for the speed of light • c = 2.998 x 108 m/s • All light waves will have the same speed, so speed is a constant

  23. Waves

  24. Wave Equations • Without the energy component c = λ • c = 2.998 x 108 m/s • λ (lambda) = wavelength in meters • (nu) = frequency in 1/s or s-1 or Hz

  25. Wave Equations • With the energy component E = h or E = hc/λ • E = energy in Joules • h =Planck’s constant= 6.636x10-34 J·s

  26. Quantum • A specific quantity of light • Bohr said that when energy is added to atoms, the electrons gain a “quantum” of energy to move to a higher level. • When electrons relax back to their normal state, they emit a quantum of energy to go back to the lowest level.

  27. Quantum…photon • Photon is just the name for a quantum of light • Electron Transition – when an electron moves from one level to another • When an electron transitions to a higher energy level, a photon is absorbed. • When an electron transitions to a lower energy level, a photon is emitted.

  28. Quantum…photon • The emitted photon is just a “piece” of light. • It has a specific energy value, so it has a specific wavelength, frequency and color • If you can measure the wavelength of the photon, you can calculate its energy.

  29. Example • The photon released by a certain electron transition has an energy of 4.56x10-19 J. Calculate the wavelength and frequency of this light. Is it in the visible range? • E=h so =E/h • (4.56x10-19 J)/(6.626x10-34J·s) • = 6.88x1014Hz • c=λ so λ=c/ • (2.998x108m/s)/(6.88x1014Hz) • = 4.36x10-7m = yes in the visible range

  30. Gas Discharge Tubes • Another way to give energy to the atom is using electricity

  31. Gives a spectrum just like that of a flame… Figure 5.12 in your textbook

  32. Colors…energy? • Once you read the wavelength from your spectroscope scale, you can calculate the energy the electrons had to lose in order to release that color of photon.

  33. Ta-daa • This is why scientists can calculate the energy values of the levels within an atom even though they can’t see them! • Please complete the packet, “Analysis of Spectral Lines”