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In this experiment students have the opportunity to notice and acknowledge how the process of spectrometry actually works and they can also observe how different substances emit different colors due to the specific wavelength that is emitted. The objectives of this experiment were accomplished through the initial spectroscope calibration and its associated calibration graph, the comprehension of how an emission spectrum works, and the knowledge that different metals give off characteristic flame emission spectrums. The general method associated with part one is to essentially just calibrate the spectroscope by looking through it against the fluorescent light as a zero reference. Part two just consists of looking through the spectroscope and observing the spectrum using the hydrogen lamp. Once the scale reading is determined, the observed wavelength can then be calculated through the use of the equation obtained from the calibration graph. In order to determine the energy transition that each of the observed lines represents, the initial n value must be found which is done through the use of the equation: E=-2.178×10-18J((1/n2final)-(1/n2initial)). Part three is primarily just associated with the burning of NaCl, CaCl2, LiCl, SrCl2, BaCl2, and KCl in order to examine the colors emitted through the use of a spectrometer and without the use of a spectrometer.
For part one the spectrometer is the main focus of the experiment. A spectrometer is an instrument that allows people to view a specific light as a set of definite wavelengths. It must be calibrated in order to make sure that the systematic error can end up being corrected. This is primarily associated with how the readings obtained from the spectroscope will not precisely correspond with the accepted values, therefore calibration of the instrument is quite important and necessary for the experiment. In order to properly calibrate the spectroscope first, obtain a set of known emission wavelengths for a gas that is available from a reference source. In this case the set of known emissions is given from the mercury lines observed in the fluorescent light. The mercury emission line spectrum has a wavelength of 404.7nm for violet, 435.8 nm for blue, 546.1 nm for green, and 579.0 for yellow. These lines represent the individual components of the light source, in this case the mercury form the fluorescent lights.After obtaining the set of data, record the wavelengths observed for the gas using the spectroscope. Prepare a graph of known (reference) wavelengths versus observed wavelengths. From this calibration plot, observed emission lines can be corrected to agree with wavelengths recorded with more accurate instrumentation. The calibration graph had a positive slope when we drew a linear least square fit. The equation of the line and R2 was
R2 = 0.9268
The y value corresponds to the spectroscope scale reading and the x value corresponds to the wavelength (nm). A source of error associated that could possibly lead to this difference is how there was some slight interference with the hydrogen lamp readings due to how there was some fluorescent light in the background that could have impacted the values we obtained, thus our calculations to find the initial n value would be swayed for each of the colors emitted.
Part two regarding the electronic transitions we obtained for red, green, blue and violet the initial n values are as follows: red(3), green(3), blue(5), and violet(8). The known transitions for the hydrogen atomic emission spectrum is as follows: red(3), green(4), blue(5), and violet(6). The values we obtained for green and violet differ by a value of one and a value of two. In regards with the wavelength values that we obtained in part two the percent error for violet is 4.57%, for blue it is 3.79%, for green it is 15.94%, and for red it is 5.15%. Green has the highest percent error and that corresponds with how green has the greatest difference in the initial n value. Also the cover over the lights could have affected our results. If this was the case it would have thrown our numbers off when we calculated the observed wavelength of the hydrogen lamp. This could be what accounted for our percentage errors in part two. Another source of error associated with part two could be that we did not properly line up the slit in the spectroscope which would give us an inaccurate scale reading.
The main concept associated with part three is that of how when atoms are heated the excited electrons go up to higher energy levels and when the electrons drop they emit a quantum of energy, the wavelength(color) of the light that is observed is the difference between the two energy levels. The only transitions observed that we can see are the ones that are associated with the visible wavelength. NaCl had an orange flame and through the spectroscope only the orange line appeared, CaCl2 had an orange red flame and through the spectroscope has a green line. The SrCl2 appeared red with and without the spectrometer while the LiCl had a violet and red flame and through the spectrometer appeared just red. Lastly, the KCl had a pale indigo flame looked green/yellow with the spectrometer and the BaCl2 had a yellow green flame and was solely green with the spectrometer. All of the metal salt’s spectrums appeared to only contain one or two colors and the colors in the spectrum did not vary much from the color seen with the naked eye. To clarify, if someone saw the color without a spectroscope, then looked through the spectroscope, it would make sense to them that the metal salt had that spectrum. Lastly, as the energy of the light decreased, so did the intensity.
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