Spectral Analysis of Metals & THE HYDROGEN SPECTRUM
PART I: METALLIC SALTS
Introduction
In the early 1900’s, scientists began utilizing their observations of matter to explain chemical behavior. For instance, they had observed that certain elements emitted visible light when heated in a flame. According to the Bohr model of the atom, electrons may occupy only specific energy levels. When an atom absorbs sufficient energy, an electron can “jump” to a higher energy level. When these excited electrons fall back to lower energy levels, they release excess energy in packages of light call photons of electromagnetic radiation, or light quanta (if the wavelength of the released photon is between 400 nm and 700 nm, the energy is emitted as visible light).
An electron may drop all the way back down to the ground state in a single step, emitting a photon in the process. Alternatively, an electron may drop back down to the ground state in a series of smaller steps, emitting a photon with each step. In either case, the energy of each emitted photon is equal to the difference in energy between the excited state and the state to which the electron relaxes. The flame color may be described in terms of its wavelength, and the following equation may be used to calculate the energy of the emitted photon:
∆E = (hc/λ)
∆E is the energy difference between the two energy levels in Joules (J)
h is Planck’s constant; h = 6.63 x 10-34 J∙s
c is the speed of light; c = 3.00 x 108 m/s
λ is the wavelength of light measured in meters; usually given in nanometers (nm) (1 m = 1.00 x 109 nm)
The color of the light depends on the specific energy change that is taking place. White light is a continuous spectrum in which all wavelengths of visible light are present. An excited atom, however, produces only one or more specific lines of light in its spectrum, corresponding to the specific changes in energy levels of its electrons. When heated, each element emits a characteristic pattern of light energies, which is useful for identifying the element. The characteristic colors of light produced when substances are heated in the flame of a gas burner are the basis of flame tests for several elements.
Representative Wavelength (nm) / Wavelength Region (nm) / Color410 / 400 – 425 / Violet (lilac)
470 / 425 – 480 / Blue
490 / 480 – 500 / Blue-green
520 / 500 – 560 / Green
565 / 560 – 580 / Yellow-green
580 / 580 – 585 / Yellow
600 / 585 – 650 / Orange
650 / 650 – 700 / Red
Purpose
You will observe the spectrum of certain elements using the flame test and then be able to explain the relationship between the color of light and the energy it represents. This energy will correspond to the falling of electrons from a higher energy level to a lower energy level.
Materials
safety gogglesspectroscopematches
Q-tips© soaked in waterplastic cupsBunsen burner
colored pencilsMetal salts: Ca+2, Cu+2, Na+, Li+, K+, Sr+2
Procedures
1. Your instructor will demonstrate how to proceed with the flame test.
2. At each lab station, you will find seven labeled plastic cups containing a small amount of a metallic salt.
3. Dip the Q-tip© into the plastic cup to pick up a few grains of the metallic salt and place in the hottest part of the burner flame. Do not leave the Q-tip© in the flame too long or it will burn.
4. Record the flame color in your data table.
5. View the resulting flame through the spectroscope. Draw the observed lines using colored pencils in your data table.
6. Repeat steps 3 and 4 for each salt.
Data
METALLIC SALT / FLAME COLOR / LINE SPECTRALiCl
NaCl
KCl
CaCl2
SrCl2
CuCl2
RESULTS
Using the representative wavelength for your observed color, complete the results table below:
Metal / Flame Color / λ (nm) / λ (m) / ∆E (J)Part II. THE HYDROGEN SPECTRUM AND THE BOHR MODEL
Introduction
In 1913 Niels Bohr published an explanation of why the hydrogen spectrum looks the way it does. The spectrum of hydrogen was well known before Bohr stepped onto the scene, but there was no good explanation of why there were individual lines, and why they had the wavelengths they did. Years earlier (1885) J. Balmer, a Swiss mathematics teacher, had developed a formula which calculated the wavelengths. Balmer took on the problem in his later life because he was interested in solving numerical puzzles. Despite publishing a paper on the hydrogen spectrum, he did not offer a physical explanation of the hydrogen spectrum.
Bohr developed a fully functional model of the hydrogen atom, based on Rutherford’s explanation of the nucleus and Planck’s Quantum Theory. Briefly, Bohr suggested that the electrons “orbited” the nucleus at only specific distances which were related to whole-number multiples of a constant. Each of these orbits had a specific energy. Bohr went on to say that electrons in these “allowed orbits” would not radiate energy (which deviated from classical physics). If an electron moved from one allowed orbit to another it either absorbed a specific amount of energy or emitted a specific amount of energy. The energy involved in moving from one orbit to another was the mathematical difference between the energies of the two orbits.
Our work with the hydrogen spectrum is based on two formulas. The first computes the amount of energy an electron has in a specific orbit. The second gives the relationship between the energy and wavelength of electromagnetic radiation and is used to compute the wavelength of the light energy emitted as an electron moves from a higher orbit to a lower orbit.
Notes:
En is the energy at a specific orbit (n = 1, 2, 3, … 7) or allowed energy state in Joules
Rh is the Rydberg Constant, with a value of 2.18 x 10-18 J
n is the orbit number (1, 2, 3 … 7); we now know n as the principal quantum number
E is the energy associated with a wavelength of light in Joules; from E = hf where c = f
h is Planck’s Constant and has a value of 6.63 x 10-34 Js
c is the speed of light and has a value of 3.00 x 108 m/s
is the wavelength of light (or any electromagnetic radiation) in meters
We also make use of the convention that E is the difference between two energies.
E = Efinal – Einitial
Gas Discharge Tubes
A bright-line spectrum is produced by hot gases of low density. The spectrum that results has bright lines separated by dark spaces. These bright lines are determined by the kinds of atoms present in the gases and the amount of energy supplied. Each gas emits its own bright-line spectrum.
Caution: The power packs for the gas tubes are high voltage. Do not touch the exposed leads to the power packs.
Purpose
You will observe the spectrum of certain elements using high voltage emission from a gas and then be able to explain the relationship between the color of light and the energy it represents. This energy will correspond to the falling of electrons to a lower energy level.
Materials
gas tubesincandescent lightfluorescent light
spectroscopescolored pencils
Procedure
1. Using the spectroscope examine and sketch, in color, the spectrum observed for the following: fluorescent light, neon, hydrogen, helium, mercury, krypton and an unknown.
Data
Use colored pencils to record the colors of the spectrum for each light. When the spectrum is continuous shade the colors together. If the spectrum is a line spectrum, use individual vertical lines of color.
LIGHT / OBSERVED SPECTRAFluorescent
Ne
H
Ar
Hg
Experimental- With a photograph of H-spectrum
1.View the photograph of the spectrum of hydrogen in the PowerPoint slideshow.
2. In your notebook record (1) the color, and (2) the wavelengths of the spectral lines you observe (to three significant digits – you will have to estimate the third digit).
Sketch your observations, and label the colors and their wavelengths in meters and nanometers.
350 nm 700 nm
Calculations
Each group will compute the wavelength of a transition in the hydrogen spectrum and put the results in a table on the white board at the front of the room. You will be assigned a transition by your teacher. Record the n values that you are given below:
n initial = ______ n final = ______
(4 points)
(1) Use the equation to compute the energies at each of the two energy levels.
E initial = ______
E final = ______
(1 point) (2) Compute the difference in energy between the two energy levels: E = Efinal – Einitial
E = ______
(4 points)
(3) Start with and solve it for , then compute the wavelength. Be sure to show the units for every number. Compute the wavelength in both meters and nanometers.
in meters = ______
in nanometers =______
(4) Put your results (just the wavelength in nanometers) in the table on the screen.
Results
Copy the values from the table on the screen into your table below. Nothing will go in the cells in the table with shading. These represent energy being absorbed and there would be no light emitted. Look for matches between the wavelengths that you and your classmates calculated and the wavelengths you measured in the hydrogen spectrum.
Hydrogen Spectrum Wavelengths by Energy Level Transition
wavelengths in nanometers
Electrons come from these higher energy levels and go to …
…these lower energy levels. / 7 / 6 / 5 / 4 / 3 / 26
5
4
3
2
1
Questions (14 points)
(2 points) 1. From the flame test describe the relationship between the colors you observed without the spectroscope and the line spectra you observed through the spectroscope.
(1 point) 2. What color of light in the electromagnetic spectrum has the highest frequency?
(1 point) 3. What color of light in the electromagnetic spectrum has the lowest frequency?
(1 point) 4. To what energy level do the electrons drop when visible light is produced?
(1 point) 5. What kind of radiation is produced when the electrons move from a higher level to the first energy level?
(1 point) 6. What kind of radiation is produced when the electrons move from a higher level to the third energy level?
(2 points) 7. Why is the energy (E) negative?
(1 points) 8. A glass rod was heated in a burner flame and gave off a bright yellow flame. What metal ion predominates in the glass rod?
(2 points) 9. The alkali metals cesium (Cs) and rubidium (Rb) were discovered based on their characteristic flame colors. Cesium is named after the sky and rubidium after the gem color. What colors of light do you think these metals give off when heated in a flame?
(2 points) 10. Aerial fireworks contain gunpowder and chemicals that produce colors. What elements would you include to produce crimson red? Yellow?
OVER