1. Light CHEM HONORS

2. The Nature of Light Light is electromagnetic radiation, a wave composed of oscillating, mutually perpendicular electric and magnetic fields propagating through space. In a vacuum, light travels at a constant speed of 3.00 x 10 8 m/s

3. Characterization of a Wave Amplitude – the vertical height of a crest ( or depth of a trough) The greater the amplitude the more intense or bright the light is Wavelength ( ) – distance between adjacent crests (commonly measured in m, μm, or nm) Waves with a large amplitude and a short wavelength are the most energetic

4. Frequency Frequency ( ) – the number of cycles that pass through a stationary point in a given period of time. The units are s -1 (cycles/s) or Hz ( 1 cycle/s) The frequency is directly proportional to the speed at which the wave is traveling and is inversely proportional to the wavelength ( ) = c /

5. Checkpoint Calculate the wavelength (in nm) of the red light emitted by a barcode scanner that has a frequency of 4.62 x 10 14 s -1. = c /

6. Electromagnetic Spectrum

7. Interference

8. Diffraction

9. Interference from Two Slits

10. The Particle Nature of Light After the discovery of the diffraction of light, light was thought to be purely a wave phenomenon. However a number of discoveries brought this classical view into question – mainly the photoelectron effect

11. The Photoelectron Effect The observation that many metals emit electrons when light shines upon them. Classical electromagnetic theory would attribute the amplitude of the wave (intensity of the beam) as the source of the emitted electron The experimental results did not support this theory – high frequency, low intensity lights were still emitting electrons The results showed that there was a threshold frequency In other words, low frequency light does not emit electrons whereas high frequency light does.

12. The Photoelectric Effect Einstein used quanta to explain the photoelectric effect. Each metal has a different energy at which it ejects electrons. At lower energy, electrons are not emitted. He concluded that energy is proportional to frequency: E = h where h is Planck’s constant, 6.626  10 −34 J ∙ s.

13. The Nature of Energy—Quanta Max Planck explained it by assuming that energy comes in packets called quanta (singular: quantum). Planck’s constant – the energy of each photon in terms of the photon’s frequency.

14. Checkpoint A nitrogen gas laser has a wavelength of 337 nm, what is the energy of the laser? E = h = c / c = 3.0 x 10 8 m/s h = 6.626 x 10 -34 J s

15. Atomic Spectroscopy and the Bohr Model Another mystery in the early twentieth century involved the emission spectra observed from energy emitted by atoms and molecules.

16. Continuous vs. Line Spectra For atoms and molecules, one does not observe a continuous spectrum (the “rainbow”), as one gets from a white light source. Only a line spectrum of discrete wavelengths is observed. Each element has a unique line spectrum.

18. The Hydrogen Spectrum Johann Balmer (1885) discovered a simple formula relating the four lines to integers. Johannes Rydberg advanced this formula.

19. Questions (Read Pages 306 – 309) 1.) What causes spectral lines? Answer on the atomic level and use Planck’s and Bohr’s ideas 2.) Why do elements have more than one spectral line? Why aren’t there infinitely many lines? Explain the significance, with respect to proving or disproving the atomic model, of there not being infinite many lines 3.) According to Bohr’s atomic model, where may an atom’s electrons be found? 4.) What does it mean for electrons to become “excited”?

20. Questions 5.) State the equation that is used to determine the energy content of a packet of light of specific frequency. If only the wavelength of light is known, what additional equation is needed? 6.) What form of energy emission accompanies the return of excited electrons to their ground state? 7.) What is the frequency of a photon of light that has a wavelength of 692 nm? What is the energy of one mole of these photons (in kJ/mol)? 8.) How does a flame test support the quantum mechanical model of the atom? Why are different metals different colors? 9.) Explain the difference between emission and absorption spectra?