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Particle Nature of Light

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  1. Particle Nature of Light • New Experiments showed that the wave model of light needed to be revised. • The wave model of light cannot explain why heated objects emit only certain frequencies of light at a given temperature • Or why some metals emit electrons when colored light of a specific frequency shines on them.

  2. Quantum Concept • Max Planck (1900) • Observed - emission of light from hot objects • Concluded - energy is emitted in small, specific amounts (quanta) • Quantum - minimum amount of energy that can be gained or lost by an atom

  3. Classical Theory Quantum Theory Quantum Concept • Planck (1900) vs. Because these steps are so small, the energy/temperature seems to rise in a continuous, rather than a stepwise, manner.

  4. The quantum concept

  5. Remember: Energy is related to frequency (and also wavelength!) • Planck’s constant has a value of 6.626 x 10–34 J · s, where J is the symbol for the joule, the SI unit of energy. • Looking at the equation, you can see that the energy of radiation increases as the radiation’s frequency, v (of “f”), increases.

  6. The Quantum Concept • According to Planck’s theory, for a given frequency, n, (or f) matter can emit or absorbenergy only in whole-number multiples of hn; that is, 1hn, 2hn, 3hn, and so on. • Matter can have only certain amounts of energy—quantities of energy between these values do not exist.

  7. Light as a Particle • Einstein (1905) • Observed - photoelectric effect

  8. The Photoelectric Effect • In the photoelectric effect, electrons, called photoelectrons, are emitted from a metal’s surface when light of a certain frequency shines on the surface.

  9. Light as a Particle • Einstein (1905) • Concluded- light has properties of both waves and particles “wave-particle duality” • Photon - particle of light that carries a quantum of energy

  10. The Photoelectric Effect • Further, Einstein proposed that the energy of a photon of light must have a certain minimum, or threshold, value to cause the ejection of a photoelectron. • According to this theory, even small numbers of photons with energy abovethe threshold value will cause the photoelectric effect.

  11. Bohr’s Model of the Atom • A combination of Plank’s theory and Einstein’s theories helped Bohr develop and explain his EnergyLevelModel. • Bohr studied the Atomic Emission Spectrum of the Hydrogen Atom to calculate the energy of each of the energy levels in the atom.

  12. Atomic Emission Spectra of Hydrogen • Hydrogen’s atomic emission spectrum consists of several individual lines of color, not a continuous range of colors (like a rainbow) as seen in the visible spectrum. • The fact that only certain colors appear in hydrogen’s atomic emission spectrum means that only certain specific frequencies of light are emitted.

  13. Atomic Emission Spectra • The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by atoms of the element. • An atomic emission spectrum is characteristic of the element being examined and can be used to identify that element.

  14. Bohr Model • Bohr proposed that electrons exist only in orbits with specific amounts of energy called energy levels • Therefore… • e- can only gain or lose certain amounts of energy • only certain photons are produced • These photons have different colors because they have different energy values

  15. Line-Emission Spectrum •Bohr said the different colors of light emitted were caused by an electron that had been excited away from their ground state near the nucleus, returning to its ground state. excited state ENERGY IN PHOTON OUT ground state

  16. The lowest allowable energy state of an electron in an atom is called its ground state. • Electrons must absorb a specific “quantum of energy” to be excited to a higher energy level. This is said to be it’s “excited state” • When the excited electron drops back down to its “ground state” it emits a photon of light with the same energy as the difference between the energy levels it drops.

  17. Energy Level Model • Bohr measured the wavelengths and frequencies of the light emitted and looked for patterns. • He came up with an equation to determine the energy of the light that will be emitted or absorbed when an electron changes energy levels. • This is how he calculated the energy of each of the Energy Levels in his model!!!

  18. Energy of photon depends on the difference in energy levels Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom Bohr Energy Level Model 6 5 4 3 2 1

  19. The smaller the electron’s orbit, the lower the atom’s energy state, or energy level. • Conversely, the larger the electron’s orbit, the higher the atom’s energy state, or energy level.

  20. An explanation of hydrogen’s line spectrum • The four electron transitions that account for visible lines in hydrogen’s atomic emission spectrum are shown.

  21. Other Elements • Each element has a unique bright-line emission spectrum. • “Atomic Fingerprint” Helium • Bohr’s calculations only worked for hydrogen! 

  22. Energy states of hydrogen • And although a hydrogen atom contains only a single electron, it is capable of having many different excited states.