Electron Emission Spectra

Overview Types of waves Parts of a wave Relating wavelength and frequency Electromagnetic spectrum Atomic spectrum Calculate the wavelength of light emitted Relating it all to Heisenberg’s Uncertainty Principle

Parts of a wave In Physical Science and Earth Science you (should have) learned about two types of waves: Longitudinal a.k.a. compression waves Sound waves and P waves from earthquakes Transmit energy as vibrations that move matter Cannot travel in a vacuum because they need to bounce off of atoms Transverse Light and water waves as well as S-waves from earthquakes Transmit energy in small particles of energy called photons The various wavelengths make up the electromagnetic spectrum

Parts of a Transverse Wave Increased amplitude Increased wavelength, decreased frequency

Frequency and Wavelength Relationship Frequency is the number of waves (or cycles) that pass a particular location in 1 second. The SI unit is Hertz (or s-1) The symbol is the greek letter gnu (ν) Wavelength is the length of a wave The SI unit for wavelength is meters (m) The symbol is the greek letter lambda (λ) Frequency and wavelength are inversely proportional As frequency increases, the wavelength decreases As frequency decreases, the wavelength increases They are related by the following equation: c=λν Where c is the speed of light (2.998 x 108 m/s) You will be expected to use this equation to determine the wavelength if given the frequency of light and vice versa. (pg 140)

The Electromagnetic Spectrum Light, as we know can have various wavelengths and frequencies. These fit an entire spectrum called the electromagnetic spectrum As you move to the right, the waves become increasingly dangerous Find the following: Waves with the longest λ Waves with the shortest λ Waves with the highest ν Waves with the lowest ν Color of light with the highest ν Color of light with the lowest ν Color of light with the longest λ Color of light with the shortest λ What is the speed of red light? What is the speed of blue light?

Now using your reference tables, answer the following questions What is the frequency of green light? What is the wavelength of blue light? What is range of frequencies people can see? What is the range of wavelengths in the visible spectrum? What is a wavelength for the infrared heat lamp used to keep your food warm at McDonald’s? Is this more or less dangerous than visible light?

Atomic Emissions We’ve learned: What we’ve not touched on yet is: about the electrons being found in various levels in an atom. electrons prefer to stay in electron shell that requires the lowest energy before filling those that need higher energies to be occupied (i.e. we fill the s-shell before the p-shell, because it requires less energy) What we’ve not touched on yet is: If we give an electron extra energy, it becomes excited and jumps. If the electron is given enough energy it can jump to a different electron shell. jump from s to p, d, or f Jump from p to d or f After a while, the electron falls back to a lower energy shell. Because it’s dropping back to a lower shell, it must give off energy The energy is released as light The photon of light varies depending on the amount of energy released. The equation E=hv can be used to determine the frequency of light emitted, (where h = Plank’s constant) but we will not use this equation this year. If you combine the various colors emitted from a single element, you create an emission spectrum

Apply it to the Bohr Model This works great for Hydrogen! If we give an electron extra energy, it becomes excited and jumps. If the electron is given enough energy it can jump to a different electron shell. jump from s to p, d, or f Jump from p to d or f After a while, the electron falls back to a lower energy shell. Because it’s dropping back to a lower shell, it must give off energy The energy is released as light The photon of light varies depending on the amount of energy released. If you combine the various colors emitted from a single element, you create an emission spectrum

The Bohr model worked great for Hydrogen! However, the emission spectra for other elements did not fit with the idea that there were only certain levels for the electrons to move to/from. The quantum mechanical model was developed to support the complex atomic emission spectra found for other elements This required there to be s, p, d, and f shells in each level

Lyman, Balmer, and Paschen Series emits ultraviolet light Fall to the n = 1 energy level Balmer emits visible light Fall to the n = 2 energy level Paschen emits infrared light Fall to the n = 3 energy level

The point…. The quantum mechanical model was developed. Light travels as a wave in a ball of energy we call a photon The electrons can get excited and give off their energy as they fall back to a lower energy shell This relates back to Heizenberg’s Uncertainty Principle