Introduction: Flame Tests: When a substance is heated in a ...



Name:

Atomic Spectrum – Flame Test Lab

Introduction: Back in the 18th century, chemists began using flame tests to identify and distinguish elements. Different elements produce different colored flames when placed in the fire of a Bunsen burner. The different colors of flames are produced by electrons moving from a higher to a lower energy level and are useful in identifying an element.

Many of the flame colors appear similar when viewed in a dark room. To get a more detailed view of the flame, a spectroscope can be used to separate the colored light into distinct colored bands. Like a prism, the spectroscope bends the light waves differently depending on their wavelength. Long wavelength red light is bent the most and short wavelength violent light is bent the least. The combination of colored lines produced by the spectroscope is called the atomic line spectra and is UNIQUE, like a fingerprint, for each element.

Neils Bohr studied the atomic line spectra for hydrogen and investigated the relationship between the line spectra and the structure of the atom. He proposed that electrons in an atom can only have specific energy values, called energy levels, and that electrons could jump up to a higher level only when they absorbed a photon of light that exactly matched their energy value. By exciting the electrons of an atom with energy from a Bunsen burner we cause electrons to jump to a higher energy level.

At some point the electrons will lose that excess energy and fall back to their lower, ground state, energy level. It is the emission of this excess energy that results in the colored bands on the atomic line spectra. Each colored line represents a transition between a particular pair of energy levels. For example, a red line might indicate a transition between energy level 1 and 2. A violet line may indicate a transition between energy levels 1 and 4. The energy of these lines can be calculated from its wavelength and thus can give us a picture of how the different energy levels are structured within the atom. Because each element is unique, its set of atomic line spectra is also unique. The wavelengths of visible light are given in units of nanometers. See Table 1.

Safety: Wear goggles and aprons. Use care when lighting and working with the Bunsen burner. Do not allow the various chemicals to come into contact with your skin.

Purpose: In this lab you will investigate the characteristic flame color and spectral lines produced by solutions of different metal ions. You will need to make careful, descriptive observations of the colors and then use those descriptions to identify an unknown sample.

Pre Lab Questions:

1. What is a flame test?

2. Write the orbital notation for the following elements:

• Copper

• Strontium

• Barium

• Lithium

• Potassium

• Calcium

3. Write the complete electron configuration notation for the following elements:

• Copper

• Strontium

• Barium

• Lithium

• Potassium

• Calcium

4. Write the abbreviated or noble gas notation for the following elements:

• Copper

• Strontium

• Barium

• Lithium

• Potassium

• Calcium

Procedures:

Part 1: 1. As a class observe the various gases as Mrs. Hodgson “excites” the gases. Observe them with both your “naked” eye and with spectroscope/glasses. Record your observations in Data Table 2

Part 2: 1. Set up a Bunsen burner and correctly adjust the flame until two blue cones are visible.

2. Using tongs, remove one a cotton swab from the distilled water and dip into the beaker containing copper (II) nitrate. Gently wave the swab through the upper cone of the Bunsen burner. DO NOT ALLOW THE SWAB TO IGNITE. Observe the color of the flame with your naked eye (and attempt to observe the flame with the spectroscope). Throw away the swab and record the color of the flame in Data Table 3.

3. Repeat step 2 for the remaining test solutions listed in Data Table 3. Some colors may be similar. Be descriptive with your observations, you will have to rely on them to identify your unknowns!

4. Repeat step 2 using the 2 unknown solutions. Make observations about the unknown solutions. Record your observations in the Data Table 4. The two unknowns are two of the solutions listed in Data Table 3, form a hypothesis regarding the identity of the unknown solutions (predict the identity of unknown #1 and #2 and write it as an If, then statement).

5. When finished clean and dry lab area. Draw a rectangle around your name on the first page of this lab.

Data and Observations:

|Table 1 - Visible Portion of Electromagnetic Spectrum |

|Representative |Wavelength |Color |

|Wavelength (nm) |Region (nm) |  |

|410 |400-425 |Violet |

|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 |

|Data Table 3 - Flame |  |

|Tests | |

|Solutions |Flame Color |

|Cu(NO3)2 |  |

|Sr(NO3)2 |  |

|Ba(NO3)2 |  |

|LiNO3 |  |

|KNO3 |  |

|Co(NO3)2 | |

|Ca(NO3)2 |  |

|Data Table 2 – Line Emission Spectrum of Gases |  |

|Name of Gas |Color of Light with Naked Eye |Predominant Colors of Emission Spectrum |

|  |  |  |

|  |  |  |

|  |  |  |

|  |  |  |

|Data Table 4 – Identifying Unknown Solutions|Observations |

|Using Flame Tests | |

|Unknown 1 |  |

|Unknown 2 |  |

|Hypothesis regarding Unknown #1: | |

|Hypothesis regarding Unknown #2: | |

Conclusion Questions and Analysis

1. Based on your observations, what are the identities of your two unknowns? Justify your answer.

Unknown 1 is ________________________ Unknown 2 is ______________________

2. Determine the approximate wavelength (using Table 1) of light emitted for the LiNO3 solution. Use dimensional analysis to convert the wavelength into meters. Show your work. (1X10-9m = 1nm)

3. Use the formula c=v( to calculate the frequency of red light. (LiNO3 should have emitted a red light so you should be able to use the wavelength you obtained from question 2.)

4. Use Planck’s equation (E=hν) to calculate the energy of one photon of red light. (Use the frequency you calculated in question 3.)

5. Determine the percent error for your calculation in problem 4 if the accepted value is E=3.1 X 10 –19 J.

6. Why do different elements have different line emission spectrums?

7. A fire truck arrives at the scene of an explosion at the Bohrilogical Chemical Company. The employees rushing from the building are calling “It’s the potassium nitrate …hurry!” As the fireman approach the building they notice green flashes of light throughout the flame. One of the firemen remembers his high school chemistry class doing a flame test experiment. Something about this scenario didn’t add up. What might he be thinking?

8. Calculate the energy (in J) of one photon of green light (use Data Table 1 to find the wavelength).

9. Circle the correct answer. Use page 92 in your book to help you. Which of the following has the

• shorter wavelength – indigo light or blue light?

• lower frequency – green or orange light?

• higher frequency – microwaves or TV waves?

• higher energy – UV or gamma rays?

• higher energy – UV or IR?

10. What is the wavelength of light with a frequency of 5.21 X 1014 Hz. (Show work)

11. Find the frequency of light with a wavelength of 7.3 X 10-7 meters. (Show work)

12. If the energy of a photon of light is 4.1 X 10 –19 J what is the frequency? (Show work)

................
................

In order to avoid copyright disputes, this page is only a partial summary.

Google Online Preview   Download