Spectroscopy of Atoms and Molecules

Transcription

1 CHEM 121L General Chemistry Laboratory Revision 2.1 Spectroscopy of Atoms and Molecules Learn about the Interaction of Photons with Atoms and Molecules. Learn about the Electronic Structure of Atoms. Learn about Spectroscopy. Learn about Beer s Law. In this laboratory exercise, we will probe the behavior of electrons within atoms using Emission and Absorbance Spectroscopy. We will first examine the photons emitted from excited atoms in various salts. Then we will observe the line spectrum of excited Hydrogen atoms and excited Helium atoms. Finally, we will leverage the photon absorbance of a solution of Al 3+ ions that have been complexed to Aluminon dye molecules to determine the concentration of Al 3+. The dependence of solution absorbance on concentration can be used to determine the amount of Aluminum in our previously synthesized Alum sample. This last determination, reported as the weight percentage Al in Alum, can then be compared with the wt% Al based on Alum s chemical formula; KAl(SO 4 ) 2 12H 2 O. The spectroscopic results will then provide additional evidence our synthetic procedure for the production of Common Alum is correct. Probing the behavior of electrons within atoms is problematic; atoms themselves are far too small to be seen and their presence must be inferred, and the electron itself is a quantum mechanical object. However, understanding the behavior of these electrons is important because this behavior determines an atom s chemistry. Thus, we must find an indirect probe of the electronic behavior of an atom. It is found that useful probes of this behavior are the photons of light which interact with an atom. If an interacting photon s energy matches that of an electronic transition within the atom, the photon can be absorbed. Conversely, an electronically excited atom can relax and emit a photon whose energy matches the atom s electronic transition. These photons are directly observable; therefore they provide us with a window on the behavior of electrons within an atom. In the realm of molecules, the experiment is much the same; electronic transitions of the molecule will allow for either absorbance or emission of photons. And, again, the energy of these photons can be directly measured, giving us insight into the behavior of the electrons in the molecule. Photons are quantum mechanical objects that exhibit a Wave-Particle Duality. The wavelength (λ) of a photon is related to its speed (c) via: c = λ ν (Eq. 1)

2 P a g e 2 where ν is its frequency. In a vacuum the speed of a photon is c = x 10 8 m/sec. Its energy (E) is then related to its frequency via: E = h ν (Eq. 2) where h represents Planck s constant; h = x Jsec. Photons can cover a wide range of energies, from very low energy Radio Waves to high energy Gamma Radiation. The Electromagnetic Spectrum of photon wavelengths is represented below: Typical electronic transitions within atoms and molecules are such that the corresponding photons have energies in the Visible and Ultraviolet regions of the spectrum. If a photon of the correct Energy impinges on an atom, such that this energy matches the energy required for an atomic quantum state transition, the photon can be absorbed: This phenomenon will result in a dark line in the rainbow of colors emerging from a sample radiated with White Light. Conversely, if an atom that has previously been excited into a higher quantum state relaxes, a photon who s Energy matches the difference in state energies can be emitted:

3 P a g e 3 This will appear as distinctly colored light emitted from the excited sample. In these cases, the Energy difference between the participating quantum states is related to the Energy of the photons via: E = E photon = hc / λ (Eq. 3) Thus, the photons absorbed or emitted by a sample of the atom in question are a direct probe of the energy difference between quantum states of that atom. Again, the behavior of electrons within molecules can be probed in pretty much the same manner. In the case of emissions from excited atoms, observing emitted photons of multiple wavelengths implies many quantum states may be involved in producing the spectrum. For example, three quantum states can produce emitted photons of three different wavelengths: Visually we see these photons as a characteristic color emitted by the sample. For instance, excited Sodium (Na) atoms emit visible photons of wavelengths nm, nm, and 589,5924nm; with the latter two so-called D-Lines being much more intense than the first. Thus,

4 P a g e 4 the yellow color we see emitted by Na atoms is dominated by these two intense D-Lines, which occur in the Yellow region of the spectrum. (This is why sodium vapor lamps have a yellow hue to them.) These photons can be separated by passing the emitted light through a dispersing element, such as a prism or diffraction grating, to produce a series of spectral lines; each line corresponding to photons of a given wavelength. Identifying which quantum states are involved in the emission of photons of a particular wavelength is usually difficult and will not, in general, be considered here. A major complication occurs if an absorbing or emitting species is in a condensed phase; a liquid or a liquid solution. If the absorbing molecule/atom is in a solution, it is surrounded by constantly jostling solvent molecules. Thus, each molecule/atom finds itself in a slightly different environment than its brothers. This causes the energy gap between the quantum states responsible for the absorbance of photons to be slightly different for each molecule/atom. This means we will have a series of very, very closely spaced absorbance lines. Practically, this means the Absorbance Spectrum will be a broad band, rather than a sharp line. This is as diagramed below:

5 P a g e 5 In this case, we usually identify the absorbance band by the wavelength of maximal absorbance, λ max. Measuring the wavelength of absorbed or emitted photons is accomplished in a spectrometer. The spectrometer will also measure the Intensity of the absorbance or emission. This intensity is then related to the number, or concentration, of the absorbing/emitting species. A UV-Vis Spectrometer consists of a White Light source, dispersive optics to separate the wavelengths of the light, a sample compartment and a detector. Light from the source passes through an Entrance Slit and is focused on the dispersive element, such as a diffraction grating or prism. This separates out the various wavelengths comprising the White Light into a rainbow of colors. The desired wavelength is selected by rotating the dispersive element such that the desired color passes through an Exit Slit, through the Sample and then onto a detector. The Transmittance of the light is then defined as:

6 P a g e 6 T = P / P o (Eq. 4) This is simply a measure of how many photons pass through the sample without being absorbed. This is then related to the Absorbance by: A = - log T (Eq. 5) As we have noted, the Absorbance will depend directly on the Concentration (c) of the absorbing species: Additionally, two other factors determine the Absorbance of a sample. The further light must travel through a sample, the so-called pathlength b, the greater is the Absorbance. The ability of the absorbing species to absorb light at the given wavelength, known as the extinction coefficient ε, a quantum mechanical effect, will also determine the Absorbance. The relationship between these influences and the Absorbance is given by the Beer-Lambert Law: A = ε b c (Eq. 6) Note the Absorbance is directly proportional to each factor. One last point regarding absorbance spectroscopy needs to be addressed; different colored solutions absorb different wavelengths of light. The color of the light absorbed is directly related to the color of the light transmitted; i.e., the color we see. The color absorbed is the Complement of the color we see: Color Absorbed Wavelength (nm) Color Observed Red Bluish-Green Orange Greenish-Blue Yellow Blue Yellowish-Green Violet Green Purple Bluish-Green Red

7 P a g e 7 Greenish-Blue Orange Blue Yellow Violet Yellowish-Green This simply means our spectrometer must be set to a wavelength corresponding to the maximal absorbance of the absorbing species. In this lab we will observe the light emitted by atoms that are excited in a hot Bunsen burner flame. Then we will observe the spectral lines in the emission spectrum of excited Hydrogen atoms and excited Helium atoms. Finally, we will measure the absorbance of a solution of Aluminum (Al 3+ ) Ion Complex. By knowing how this absorbance depends on concentration, we can determine the Al 3+ concentration of the solution. In our first experiment, we will simply observe the colors of the photons emitted when select atoms are excited. We will do this by placing a drop of a solution of the salt of the desired atom into a Bunsen burner flame. The heat of the flame will be sufficient to excite the atoms in question. We can then observe the light emitted when the atoms relax. We will be examining emissions from the following atoms: Li, Na, K, Ca, Sr, Ba, Cu by testing solutions of the following salts: LiCl(aq) NaCl(aq) KCl(aq) CaCl 2 (aq) Sr(NO 3 ) 2 (aq) Ba(NO 3 ) 2 (aq)

8 P a g e 8 Cu(NO 3 ) 2 (aq) Note, in each case we are relying on the fact that neither Cl - or NO 3 - will influence the emission from the cations. In our next experiment we will observe the photons emitted by excited Hydrogen atoms. In this case we will excite the atoms by placing gaseous Hydrogen (H 2 ) into a discharge tube and applying a high voltage to the sample. In the resulting electric field, the H 2 molecules dissociate and leave the resulting H atoms in an excited state (H*). H 2 2 H* (Eq. 7) H* H + photon (Eq. 8) These excited atoms will subsequently relax, giving off photons corresponding to the energy differences between associated quantum states. We will examine the resulting glow by passing this light through a prism to disperse the photons according to their wavelength. We will view the resulting spectral lines thru a spectroscope that includes a slit for focusing the light from the discharge tube, a prism and a rotatable viewer to find the spectral lines. The first model of the Hydrogen atom to explain these spectral lines was put forth by Neils Bohr in In this model, the electron orbits the nucleus in quantized orbits. In its most relaxed or Ground State, the electron orbits very close to the nucleus. When excited, the electron moves into a higher energy orbit farther from the nucleus. When it relaxes, the electron moves back to an orbit closer to the nucleus, emitting a photon in the process.

9 P a g e 9 Unfortunately this model, although predictive for the Hydrogen atom, has some physically unappealing features and is not extendable to the electronic behavior of atoms of other elements. Even Helium, which has only one additional electron, has a much more complex electronic spectrum. To illustrate this, we will observe the visible lines in Helium s discharge spectrum. In the last part of the experiment we will complex Aluminum Ion (Al 3+ ) with a dye called Aluminon (Alu) to form a reddish colored complex. Al 3+ itself does not absorb radiation in the visible part of the spectrum. Al 3+ (aq) + 3 Alu(aq) (Alu) 3 Al 3+ (aq) Colorless Light Orange Deep Reddish Orange (Aluminon Dye Molecule) (Absorbance Spectrum of Al 3+ -Aluminon Dye Complex) As can be seen in the spectrum above, the complex does have an absorbance peak at around 525 nm.

10 P a g e 10 After calibrating the concentration dependence of the absorbance of the complex, using a visible absorbance spectrometer (Shimadzu UV-255), we will dissolve some of our Alum crystals in Water and treat the solution with Aluminon. Measuring the resulting absorbance, again with the spectrometer, will allow us to determine the amount of Al 3+ in our Alum. This result should compare favorably with that of the chemical formula.

11 P a g e 11 Pre-Lab Questions 1. The D-Lines (~589nm) in the emission spectrum of the Sodium (Na) atom are due to an electronic transition from the 3p to the 3s orbital. What is the energy gap between these two orbitals of the Na atom? 2. In our flame tests we will examine the emitted light from a series of metal atoms. In each case, we are actually using a solution of the salt of these atoms. So, Chloride Ions (Cl - ) or Nitrate Ions (NO 3 - ) are also present in the solution. In assigning the spectral emissions to the metal atoms, we make some assumptions. What are two possible assumptions being made? 3. Bi 3+ complexes with Thiourea to form a colored Complex which Absorbs light at λ max = 470 nm with an Extinction Coefficient of ε = 9320 M -1 cm -1. What is the Concentration of the Complex in a solution that has a measured Absorbance A = in 1.00 cm cell.

12 P a g e 12 Procedure Flame Tests 1. Obtain a looped nichrome wire, a Bunsen burner, and the series of salt solutions to be tested. LiCl(aq) NaCl(aq) KCl(aq) CaCl 2 (aq) Sr(NO 3 ) 2 (aq) Ba(NO 3 ) 2 (aq) Cu(NO 3 ) 2 (aq) Also obtain some 6M HCl for cleaning the nichrome wire. 2. Light the burner. Dip the nichrome wire into the 6M HCl, then place the loop into the hottest part of the burner flame. Repeat this step two or three times, until the flame around the wire is nearly colorless. 3. Dip the loop of the wire into one of the salt solutions; saving the NaCl solution for last. Place a drop of the solution into the burner flame. Record your observations. Clean the wire and repeat the test for each salt solution. (When performing the test on Barium, it may be necessary to view the flame through a Cobalt Blue filter. This salt is frequently contaminated with Sodium. Thus, the Sodium emissions must be filtered out.) 4. Clean the nichrome wire and return it. Hydrogen Emission Spectrum 1. Go to the darkened room where your instructor has set-up the Hydrogen Discharge Tube. 2. Turn on the power supply and observe the emitted light through the spectroscope. You should see a series of sharp, bright spectral lines. Consult with your instructor to confirm you are indeed seeing the correct lines. 3. Record your observations. 4. View the Helium spectrum through a spectroscope. Record your observations.

13 P a g e 13 Absorbance Measurements Construction of the Calibration 1. Weigh out 0.1g of AlCl 3 hexahydrate using the analytical balance. Transfer this to a 100 ml volumetric flask. Add about 10 ml of distilled Water and swirl to dissolve the salt. If needed add some more Water, but the salt must dissolve completely. Then bring the solution to the mark with Water. Stopper the flask and invert several times to mix thoroughly. 2. Pipet, using a volumetric piper, 3.00 ml of your solution into a 25 ml volumetric flask. Dilute to the mark and mix thoroughly. This is your stock solution. 3. Prepare 50 ml dilutions of each of the following: 0 ml stock 1 ml stock 2 ml stock 3 ml stock 5 ml stock Do this in a 50 ml volumetric flask according to: Add 20 ml (meas. Grad Cyl.) Acetate Buffer [Regulates the ph to be slightly acidic. Needed for the Aluminon to be effective.] Add the Required Volume Al 3+ Stock Add 5 ml Aluminon Reagent (5 ml Vol. Pipet) Bring to the Mark with Water and Mix Allow to Stand 10 min. 4. Fill five Spectrometer Cuvettes to near the top with each respective solution. Wait until just before you are to make the absorbance measurements to do this. 5. Prepare your Alum sample before taking any spectral readings. Preparation of the Alum 1. Obtain a sample of Alum that is in excess of 0.3g. Grind this sample to a powder with a Mortar and Pestle. 2. Weigh out 0.2g of your powder on an analytical balance and place it in a small beaker. Add 15 ml distilled Water, cover the beaker with a watch glass, and heat it on a hot plate to dissolve the Alum.

14 P a g e Once the Alum is dissolved, remove the solution from the heat and allow it to cool. Transfer it to a 100 ml volumetric flask and dilute to the mark. Mix thoroughly. 4. As before, transfer 3 ml of this solution to a 25 ml volumetric and dilute it. 5. Prepare a sample for the spectroscopic measurement as before: 20 ml Acetate Buffer, 3 ml Alum solution, 5 ml Aluminon Reagent and dilute to 50 ml. 6. Fill a sixth Spectrometer Cuvette to near the top with the solution and make the absorbance measurement. Wait until just before you are to make the absorbance measurements to do this. 7. Now, proceed to the Instrument Room with your cuvettes to measure the Absorbance of the Complex for each prepared solution at 525 nm.

15 P a g e 15 Data Analysis Flame Tests No Analysis Needed Hydrogen Emission Spectrum 1. Examine the complete emission spectrum for Hydrogen provided in the Appendix. Determine the wavelengths of the emission lines you observed. (These should all be Balmer Lines.) 2. Calculate the photon energies for each spectral line you observed. 3. Prepare an Energy Level Diagram of quantum states which are responsible for the emission lines you observed. For this diagram, assume all the Balmer Lines terminate in the same quantum state. (This state is designated with quantum number n = 2.) Provide quantum number designations (n = 3, 4, 5,...) for your quantum states in the Energy Level Diagram. 4. Identify the orbits in the Bohr Model of the Hydrogen atom responsible for each quantum state. The Bohr Model provides that the radius of the electron s orbit is given by: r = x n 2 [Angstroms] (Eq. 10) where n is the state s quantum number. Calculate the radius of each of these orbits. 5. Wavelengths for all the visible spectral lines for Helium can be found at: Assign wavelengths to all the spectral lines you observed for Helium. 6. The following spectral lines are due to the indicated transitions: Line Transition nm 3d 2p nm 4d 2p What is the wavelength of the spectral line corresponding to a 4d 3d transition? In what region of the spectrum does this spectral line occur?

16 P a g e 16 Absorbance Measurements 1. Prepare a Calibration Curve from your Absorbance Measurements (Abs. vs. Concentration Al 3+ ). Use Excel to prepare the graph and add an appropriate trendline. In calculating [Al 3+ ], recall how you prepared your standard solutions: weighed out ~0.1g of AlCl 3 6H 2 O. prepared ml solution. diluted 3.00 ml of this sol'n to 2.00 ml. diluted x ml of this solution to ml Your calculation must take into account each of these steps. 2. Use your Absorbance Measurement for the Alum solution, and the trendline from your Calibration Curve, to determine the concentration of Al 3+ in the Alum solution. 3. Use this concentration determination to determine the number of moles, and then the number of grams, Al in your original Alum sample. (Review all the steps required to prepare this sample.) 4. Calculate the Weight Percentage Al in your Alum sample. 5. Compare this result with that expected based on the chemical formula.

17 P a g e 17 Post Lab Questions 1. Examine the emission spectrum of the Hydrogen atom provided in the Appendix. The first two Lyman Lines are due to n = 2 to 1 and n = 3 to 1 quantum state transitions. Use the photon wavelengths of these lines to calculate the energy differences for the associated states. Do the same for the first Balmer Line, which is due to a n = 3 to 2 quantum state transition. Are the results internally consistent? Explain. A simple energy Level Diagram may be useful. 2. In your spectrophotometer Absorbance measurements, how do you expect the errors in subsequent measurements to compare to the error in the measurement for the stock solution? 3. Identify the Oxidation State for all the metal cations used in the Flame Tests.

18 Appendix - Complete Emission Spectrum of Hydrogen P a g e 18