1. “Atomic Fingerprints” – Experiment 6

2. Atom Lab Background
Atoms are fundamental to the building blocks of matter. We are matter made of billions and billions of atoms. Atoms have three basic parts known as protons, neutrons, and electrons. The protons and neutrons are inside the nucleus of the atom and electrons circulate in the orbits of the atoms just like the planets circulate in the orbits of the sun. The focus of this lab is the electrons. The electrons exist in energy states within the atom (called orbitals by Chemists). Generally, the further away from the nucleus these states are, the higher the potential energy of the electron in that state. When an electron is “free” of the nucleus’ influence it is said to have zero potential energy – a zero energy state. When it is closer to the nucleus it has less potential energy (needs energy input to escape the influence). It is therefore at a negative energy level. When an electron jumps between energy levels of an atom it either absorbs or emits a photon. The size of that photon’s energy depends on the energy gap between the levels. This is linked to its frequency by the equation E = hf . E in this equation is energy of the photon absorbed or emitted (and the energy interval between the energy levels in joule, to convert to eV divide by 1.6 x 10 to -19). Planck constant is h and f is the frequency of the photon of electromagnetic radiation (linked to its wavelength by c = f lamda). With this background we will now proceed to do the lab.

3. Atomic Fingerprints
Learning Objectives To distinguish between a continuous spectrum and a line spectrum To compare observed and calculated lines in the hydrogen spectrum. To identify unknown elements in a fluorescent light by “atomic fingerprints” To become proficient in using a hand spectroscope

4. The Radiant Energy Spectrum
A visible spectrum is a narrow window in a broad band of radiant energy Fig. 5.8 and 5.9 on pages 120 and 121 on this slide

5. Continuous versus quantized energy

6. Continuous versus quantized energy

7. Emission spectra as “atomic fingerprints”

8. Rydberg Equation
The Rydberg Equation accounts for electrons dropping to any energy level.
Sound familiar?
These quanta of light are photons
The relationship between energy and frequency is

9. Balmer Formula – Example Exercise 1
Calculate the wavelength of light corresponding to the energy released when an electron drops from n = 6 to n = 2 in a hydrogen atom. l/L = l/91 nm (l/2 sq –l/n sq) = l/91 nm (l/2 sq. -1/6 sq.) = l/91 nm (1/4 – l/36) = l/91 nm ( ) = 0.22/91 nm After taking the reciprocal and rounding to two significant digits, we obtain L = 91 nm/ nm

10. Rydberg Equation – Example Exercise 2
Calculate the wavelength of light released when an electron drops from n = 4 to n = 1 In a hydrogen atom. 1/L = 1/91 nm (1/nL sq. 1/nH sq.) = 1/91 nm (1/1 sq. – ¼ sq.) = 1/91 nm (1/1 – 1/16) = 1/91 nm (1-0.06) = 0.94/91 nm After taking the reciprocal and rounding to two significant digits, we obtain L = 91 nm/0.94 = 96.8 = 97 nm

11. Lab Procedure Continuous Spectrum – White Light
With the hand spectroscope, observe the emission spectrum from one or more of
following: light from an overhead projector, light from an incandescent light bulb ,
and sunlight. Draw the observed spectrum on the wavelength scale in the Data Table.

12. Lab procedure B. Line Spectrum – Hydrogen
Using a hand spectroscope, observe the hydrogen spectrum from a gas discharge
tube. Draw the position of each line on the wavelength scale in the Data Table.
2. From the Balmer formula, find the wavelength (L) of light produced when electrons
drop from 3rd to 2nd energy level. (Round the answer to two significant digits; for example, nm round to 660 nm.)

13. Lab procedure
B. Line Spectrum – Hydrogen 3. Repeat the wavelength calculation for the spectral lines produced when electrons drop from 4th to 2nd energy level; and from the 5th to 2nd energy level. 4. Record the observed and calculated wavelength values in the Data Table. State the error after comparing the observed and calculated wavelengths.

14. Lab procedure
C. Line Spectra – Helium, Neon, Argon, Krypton, and Mercury
Insert a helium gas discharge tube into a spectral tube power supply, Using a hand spectroscope, observe the emission lines from helium gas. Draw the position of six intense lines on the wavelength scale in the Data Table.
2. Repeat the procedure using a neon gas discharge tube, argon gas discharge tube,
Krypton gas discharge tube, and mercury vapor discharge tube in the power supply.
Draw the position of the most intense lines on the wavelength scale in the Data Table.

15. Lab procedure D. Identifying Unknown Elements in a Fluorescent Light
Observe the line spectrum from a fluorescent light using the hand spectroscope. Disregard the continuous rainbow, and draw the position of each line on the wavelength scale in the Data Table. Compare the line spectrum from the fluorescent light to the lines in the emission spectra of He, Ar, Kr, and Hg. Identify the elements in the fluorescent light from their “atomic fingerprints.”

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