Atomic Emission Spectra Activity Spectrum of a Single Electron Element - Hydrogen • Record the line color and its positi

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Atomic Emission Spectra Activity Spectrum of a Single Electron Element - Hydrogen • Record the line color and its position for the 3 or 4 brightest lines observed using the hydrogen lamp in table 1 and calculate each wavelength using Equation 1 from your lab manual. Be sure to include a sample of your work for one of the calculations. Table 1: Bright Line Spectra for Elemental Hydrogen. Line Color Spectral Line Position (nm) 410.22 Wavelength (nm) Corrected using calibration factor purple violet 433.87 Blue-green 485.88 red 656.12
Data Analysis 1) For the first four electronic transitions in the Balmer Series, calculate the change in energy of the electron (A4), the predicted energy of the emitted photon (Eww and the predicted wavelength of the emitted photon (eta). Put the calculated values in Table 2 and be sure to clearly show an example of each calculation in the space provided. Table 2. Calculated Values for the Balmer Series of Hydrogen. ΔΕ (1) Esterol) Ashton (nm) Electronic Transition n3 → m2 n42 ns2 non2 Clearly show the following calculations for the na → na transition -- the change in energy of the electron, AE - the predicted wavelength of the emitted photon, Nonton 2) Based on your theoretical calculations, match the electronic transitions in the Balmer Series to the spectral lines you observed, and document your choices in Table 3. Then calculate the percent error between your experimentally determined and calculated wavelengths. Table 3. Comparison of Experimental and Accepted Wavelengths from the Balmer Series. Spectral Line Experimental (nm) Accepted X (nm) from Electronic Color Observed from Table 1 Table 2 Transition n3n2 Percent Error 0402 ns2 non2 Below, clearly show your percent error calculation for the na→ n2 transition.
As an example, the emission spectrum of hydrogen consists of only four visible lines: red, blue- green, violet and deep violet (although your eyes may not be sensitive to the last one, since individual's visual acuity varies). Each color corresponds to the transition of an electron from an excited state, a higher principal energy level, to a lower principal energy level, possibly the ground state or some other allowed energy level in between. As the electron drops, it emits energy in the form of a photon which may or may not be in the visible region. In this experiment you will use a spectroscope and gas discharge lamps to measure the wavelength of each bright line in the visible atomic emission spectra of both hydrogen and mercury. You will then use these measurements to calculate the photon energy for each bright line. Recall that the energy of light is related to the wavelength and frequency of the light, through Planck's constant h = 6.626 X10-34 J sec and the speed of light in a vacuum c = 3.00 x 108 m/sec. Other Equations That Might Be Useful AE = -(2.178 x 10-18 )) (1/nļfinal - 1/n_initial) (Rydberg equation, for Balmer series, Nfinas = 2] A E = hc/A [Links wavelength and energy - be careful of the units!] 1 nm = 1x109 m The final results of this experiment will be the wavelength (in meters) and the photon energy of each bright line measured for hydrogen and mercury.