Observing bright line spectra lab answers

temple university physics

Atomic Spectra

Isaac Newton and many other scientists saw that when they directed light through a prism, the light was spread out into a continuous spectrum of color. Close observation of the spectrum uncovered a puzzling pattern of fine dark lines among the bright colors. These lines were consistently in the same position when measured by different laboratories and methods, but what caused them?

As methods improved, scientists generated light from single-element sources. For example, when a pure powder of magnesium was placed in a hot flame, they saw a pattern of colored lines that was unique to magnesium. This is how the science of spectroscopy was born.

The earlier experiments by Newton (the dark lines on a continuous background) are now called absorption spectra because the dark lines appear when certain wavelengths of light are absorbed by the sample. The later experiments with pure element light sources that made bright colored lines are called emission spectra because those sources are emitting light of specific wavelengths.

Spectroscopy is simply the study of the position of the spectral lines unique to an element or molecule, and is very useful for identifying unknown molecules as well as studying intramolecular bonds. This is how astronomer Pierre Janssen discovered helium in 1868 by observing the spectrum of the sun.

Learning Goals for This Laboratory:

· Understand when the Balmer formula is applicable and how to use it calculate wavelengths

· Understand how absorption and emission spectra are measured

· Understand how electron transitions cause the observed spectra

Apparatus
discharge lamp, bulbs containing and H, Hg, Na vapor, gratings, single filament lamp, meter stick, stand

Part I. Emission Spectra
In this lab, we will first calculate the wavelengths of the hydrogen emission lines predicted by the Balmer formula, and then compare this to experiment by directly measuring the actual wavelengths of the lines.

The wavelength of the light that makes the spectral lines for hydrogen can be described by the Balmer formula

(1)

where is an integer larger than 2 and the constant = 364.5 nm.

a) Calculate the wavelengths that the Balmer formula predicts for n= 3, 4, 5, 6, 7.

Question 1. How many of these wavelengths correspond to visible light?

b) The wavelengths of the lines are found by passing the light from a hydrogen lamp through a diffraction grating which spreads out the light according to wavelength following the equation

(2)

where is the wavelength of the light for that line, is the spacing between slits in the grating, and is the angle from the central line (see Fig. 1).

The diagram below shows a top-down view of the setup for this experiment, where the grating is placed at a distance from the lamp, and spectral lines appear at the angle . A meter stick in front of the lamp allows us to measure the distance , which is used to calculate from the tangent formula.

discharge tube lamp

apparent position of spectrum line

meter stick

grating

Figure 1. Top view of experiment.

c) The video below shows the setup used for this experiment. Skip to the 4:35 mark in the video to see the spectrum of hydrogen.

https://www.youtube.com/watch?v=y84j5Bpzph4

d) Shown in Figure 2 is a still shot of the spectrum with a ruler placed just in front of the source so that the source is at the 50 cm mark. Note that 4 lines of the hydrogen spectrum are visible: red, aqua, blue and a very faint violet line.

Use the ruler to estimate the position for each of the 4 lines to the nearest millimeter. Record these values in Excel. The grating used has 600 lines/mm. Use this value to obtain the value for , the distance between consecutive lines of the grating.

e) Use the diffraction equation (Eqn. 2) in Excel to calculate the wavelengths of the hydrogen lines and compare the values to Balmer’s predicted values by calculating percent difference. The distance between the horizontal meterstick and the grating is = 73 cm.

Figure 2. Still photo of spectrum and meter stick.

Question 2. Do your experimental values generally agree with the theoretical values? Support your answer quantitatively.

Part II. Emission Spectra from Other Sources
Each line of an emission spectrum corresponds to a specific transition made by an electron between the energy levels in the atom. This is why the transitions always occur at the same energy (and wavelength) for a given atom (see Fig. 3). Refer to your text for the relationship between energy and wavelength of light.

This schematic illustrates the allowed transitions of electrons in orbitals around a hydrogen nucleus. Each ring represents an energy level for a transition. When an electron falls from a higher energy level to a lower one, it emits a photon (light). The light for a particular transition always has the same energy and therefore the same wavelength. For example, in the Balmer series shown here, an electron that relaxes from the n=3 state to a lower energy n=2 state will emit a red photon (656 nm). This transition always emits a red photon because the energy determines the wavelength (color).

n=2

3

4

5

6

7

n=1

Figure 3.

a) Watch this video showing the spectra of several elements.

https://www.youtube.com/watch?v=7_2Wi646M0o

b) Observe the mercury spectrum in the video. Record your observations on how the number of lines and their color differ from that of the hydrogen spectrum you saw in Part I.

Question 3. Why might the mercury spectrum have more lines than the hydrogen spectrum? Support your answer using what we know about where the lines come from.

Question 4. For these emission spectra, why are some lines more intense than other lines? Use the physical definition of intensity to support your answer.

Part III. Absorption Spectra
The same electronic transitions that emit light, can also absorb light. For this reason, when white light passes through a sodium vapor, for example, there will be specific wavelengths of light missing lines from the full spectrum leaving dark lines. See the video for an example of this effect with sodium.

https://www.youtube.com/watch?v=7u3rRy97m9Y

Question 5. What color of light do sodium atoms absorb?

a) Sketch a top-down diagram of the experimental apparatus used in this video that includes 4 components: the white light source, the sample (sodium vapor), the diffraction grating, and the screen.

Question 6. The white light source in this demonstration emits light because the electrons in its filament get so hot that they vibrate rapidly enough to emit radiation in the visible range. Briefly explain why this produces light that is so different from the light of the hydrogen lamp.

Question 7. The emission spectrum for sodium is shown below. Because the absorption and emission spectra mirror each other, we can see from this emission spectrum that there should be several absorption lines clustered near each other in the yellow-orange range. However, we do not see them in the video. What could be changed about the experimental setup used in the video to aid in resolving such nearby lines?

Spectroscopy in Astronomy | Astronomy