Lab Report 1 - Spectroscopy

Spectroscopy lab Introduction In this experiment, we looked at various electronic transformations in the hydrogen atom as well as other elements. Using the Bohr model of atomic structure, certain number orbits corresponding with the observed spectrum for the hydrogen atom were determined. The study of the interaction of electromagnetic radiation absorbed and emitted by atoms is known as spectroscopy. Different forms of energy changes inside a molecule can be perceived depending on the radiation. Microwave radiation for example, causes a rotation of molecules along the bond axes, which contributes information about bond lengths, whereas infrared radiation frequently causes extending and curving of bonds in a molecule. Bohr’s model, in which the electron is supposed to rotate around the nucleus in a fixed orbit, is required in order to understand why these transitions occur. In his idea, l electrons travel in circular orbits around the nucleus. electromagnetic radiation is used to irradiate the irradiating the electron in the hydrogen atom. This model argues that the electron can reside in several set distances from the nucleus. The energy can be used by the electron to move to a higher-energy orbit. Because the electron absorbs energy, it is depicted as absorption. When energy in the form of heat, light or electricity is absorbed by atoms, it often re-emits the energy in the form of light. Mercury atoms for example, produce blue light, helium atoms produce violet light and hydrogen atoms provide reddish light. An emission spectrum is the result of bright lines. A particular elements emission spectrum is always the same. Difference in energy between two stationary states is the photon energy emitted when an electron transitions from one stationary state to an alternate. Bright lines appearing on a dark background make up an emission spectrum, while dark lines on a bright background make up an absorption spectrum. When it comes to emission, an electron moves from a higher energy state to a lower level. In the middle of the same two energy levels, there is a shift in absorption, but this time from a lower to the higher one. When the white light spectrum is continuous, there is a discernible difference. This indicates that there are no abrupt changes in light intensity as a purpose of wavelength. However, the emission spectra of hydrogen, helium, and barium are not all the same. Different sorts of energy conversions within a molecule can be perceived depending on the radiation. The hydrogen atom’s lone electron is generally located in the ground state orbit. This is the configuration with the least energy and highest stability. This single electron can absorb some quanta of energy and occur stimulated to a higher orbit number as a result. When the electron reaches this higher orbit, it quickly releases the energy and returns to a more stable condition, which is the cause for the emission lines seen through the spectrometer. These emission lines are the consequence of the colors in the line spectrum, according to the conservation of energy. The principle underlying a fireworks display is this technique. The various colors are caused by elements absorbing energy in order to boost electrons to higher orbits. The energy is released as visible light when the electrons relax down to their more stable lower orbits. Another example is boiling saltwater over a flame on the stove. The flame will turn a bright yellow color which matches the line spectrum of the Na atom in the salt NaCl. We measured the wavelength of the lines corresponding to these transitions in this experiment and used that information to come to a judgement about the transition taking place. If the electron returns to n=1, the energy is released as ultraviolet light, which is invisible to our eyes. However, some of the electrons will revert to the n=2 state, and the energy released corresponds to visible light that we observed with a spectrometer. Compared to figure 2.22 in the textbook, the wavelength I got for red and blue on the hydrogen spectrum were shorter than the wavelengths in the given figure. in figure 2.22 of textbook, the green wavelength was shorter than the one I saw through the spectrometer. I noticed that the wavelengths in the given figure were more specific lengths and the ones I saw in the lab were more broad numbers. Another aspect that was different between the two hydrogen spectrogram models was the lengths. What I examined through the spectrometer was that red was the longest wavelength. In the figure, presented red was the shortest wavelength emitted from n=3 to n=2. Colum A includes each color that was shown through the spectrometer. Column B is the wavelength with the measurements I found through the spectrometer in nm. In column C, I calculated the frequency of each element, I used the equation [3.00*10^17(nm/sec)]/wavelength(nm) in hertz and just plugged in the wavelengths of each light color. Lastly, in column D I calculated the Photon using the equation (6.26*10^34 joules-sec) *v. Conclusion I observed from the experiment the spectra corresponding to electronic transitions in the hydrogen atom, mercury atom and additional sources. My experimental findings found that, red should be the shortest wavelength and blue or violet should be the bigger wavelength. I did not get those results in some of my emission lines. An error that I made to cause this was the way I looked at the measurements through the spectrometer. I should have looked at the light more through the right side of the spectrometer, but I held my spectrometer closer to the center of each light source. This caused my lines to not add up exactly from the starting measurement. For example, when I was looking through the sunlight, I got a wavelength of between 700-750 nm for violet. That was very high for this wavelength. Looking through the spectrometer with the sunlight was the easiest to see each wavelength because it was the brightest light. From this lab I learned that the emission lines visualized through the spectrometer are caused from the single electron absorbing a quantum of energy and to a higher orbit number. When it reaches this higher orbit, it quickly releases that energy and returns to its increased stable state causing the colored lines seen through the spectrometer.